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chemguide: helping you to understand Chemistry - Main Menu
chemguide Helping you to understand Chemistry You can search Chemguide either using a keyword search (courtesy of Google) or the menu system.
See the bottom of this page for the latest additions and updates.
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Notes: If you don't find what you want, use the BACK button on your browser to return to this page. Remember that this is a UK site, so spellings will be English. Take care if you want to search for words like sulphur, sulphate, sulphuric acid or aluminium.
Searching using the menus
Atomic Structure and Bonding
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Covers basic atomic properties (electronic structures, ionisation energies, electron affinities, atomic and ionic radii), bonding (including intermolecular bonding) and structures (ionic, molecular, giant covalent and metallic).
chemguide: helping you to understand Chemistry - Main Menu
Inorganic Chemistry
Includes essential ideas about redox reactions, and covers the trends in Groups 2, 4 and 7 of the Periodic Table. Plus: lengthy sections on the chemistry of some important complex ions, and of common transition metals.
Physical Chemistry
Covers rates of reaction including catalysis, an introduction to chemical equilibria, redox equilibria, acid-base equilibria (pH, buffer solutions, indicators, etc) and phase equilibria (including Raoult's Law and the use of various phase diagrams).
Instrumental analysis
Explains how you can analyse substances using machines - mass spectrometry, infra-red spectroscopy and NMR.
Basic Organic Chemistry
Includes help on bonding, naming and isomerism, and a discussion of organic acids and bases.
Properties of organic compounds
Covers the physical and chemical properties of compounds on UK A level chemistry syllabuses.
Organic Reaction Mechanisms
Covers all the mechanisms required by the current UK A level chemistry syllabuses.
About this site
Includes a contact address if you have found any difficulties with the site.
Chemistry Calculations
A description of the author's book on calculations at UK A level chemistry standard.
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chemguide: helping you to understand Chemistry - Main Menu
Textbook suggestions
Suggestions for textbooks and revision guides covering the UK AS and A level chemistry syllabuses, with links to Amazon.co.uk if you want to follow them up.
Download syllabuses
For UK A level students. Download a copy of your current syllabus from your Exam Board.
Test yourself
A link to Dr Phil Brown's website where UK GCSE, AS and A level chemistry students will find a wide and growing range of multiple choice, short answer and structured questions.
An examiner's view
A link to Rod Beavon's chemistry pages. Rod Beavon is chief examiner for A level chemistry for the UK exam board Edexcel. A close look at what he has to offer is a must for Edexcel students, but there is a lot of good stuff whatever exam system you are working in.
Latest additions and important updates
5/12/2004
There is now the beginnings of a section on phase equilibria. I'm working on it at the moment and it is unlikely to be finished before the end of January. It currently deals with vapour pressure, phase diagrams for pure substances and for solutions of non-volatile solutes (including the effect of the solute on the boiling point and freezing point of the solvent).
18/12/2004
An introduction to phase diagrams involving eutectic mixtures is now available using the tin-lead system.
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chemguide: helping you to understand Chemistry - Main Menu
© Jim Clark 2004
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Understanding Chemistry - Atomic Structure and Bonding Menu
Understanding Chemistry
ATOMIC STRUCTURE AND BONDING MENU
Basic atomic properties . . . Includes a discussion of orbitals, electronic structures of atoms and ions, ionisation energies, electron affinities, atomic and ionic radii. Bonding . . . Includes ionic, covalent, co-ordinate (dative covalent) and metallic bonding as well as intermolecular attractions like Van der Waals forces and hydrogen bonding. Also includes full discussions of electronegativity and shapes of molecules and ions. Types of structure . . . Describes and explains how the various types of structure (ionic, giant covalent, metallic, and molecular) affect physical properties.
Go to Main Menu . . .
© Jim Clark 2000
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Atomic properties menu
Understanding Chemistry
ATOMIC PROPERTIES MENU
Simple background . . . Revises the simple knowledge you should already have about the structure of atoms from introductory courses (e.g. GCSE). Atomic orbitals . . . Explains what atomic orbitals are and discusses their shapes and relative energies. This is essential pre-reading before you go on to any of the remaining topics in this section. Electronic structures . . . How to work out and write the electronic structures for atoms and simple monatomic ions (containing only one atom - e.g. Cl- or Mg2 +) using s, p, d notation. Ionisation energies . . . Explains what ionisation energies are and how and why they vary around the Periodic Table. Electron affinities . . . Explains what electron affinities are and how and why they vary around the Periodic Table. Atomic and ionic radii . . .
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Atomic properties menu
Looks at the various measures of atomic radius, and explains how and why atomic radii vary around the Periodic Table. Also considers how the radii of positive and negative ions differ from the atoms they come from.
Go to atomic structure and bonding menu . . . Go to Main Menu . . .
© Jim Clark 2000
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a simple view of atomic structure
A SIMPLE VIEW OF ATOMIC STRUCTURE
This page revises the simple ideas about atomic structure that you will have come across in an introductory chemistry course (for example, GCSE). You need to be confident about this before you go on to the more difficult ideas about the atom which under-pin A'level chemistry.
The sub-atomic particles Protons, neutrons and electrons.
proton neutron electron
relative mass 1 1 1/1836
relative charge +1 0 -1
Beyond A'level: Protons and neutrons don't in fact have exactly the same mass - neither of them has a mass of exactly 1 on the carbon-12 scale (the scale on which the relative masses of atoms are measured). On the carbon-12 scale, a proton has a mass of 1.0073, and a neutron a mass of 1.0087.
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a simple view of atomic structure
The nucleus The nucleus is at the centre of the atom and contains the protons and neutrons. Protons and neutrons are collectively known as nucleons. Virtually all the mass of the atom is concentrated in the nucleus, because the electrons weigh so little.
Working out the numbers of protons and neutrons No of protons = ATOMIC NUMBER of the atom The atomic number is also given the more descriptive name of proton number. No of protons + no of neutrons = MASS NUMBER of the atom The mass number is also called the nucleon number.
This information can be given simply in the form:
How many protons and neutrons has this atom got? The atomic number counts the number of protons (9); the mass number counts protons + neutrons (19). If there are 9 protons, there must be 10 neutrons for the total to add up to 19.
The atomic number is tied to the position of the element in the Periodic Table and therefore the number of protons defines what sort of element http://www.chemguide.co.uk/atoms/properties/gcse.html (2 of 7)30/12/2004 11:00:08
a simple view of atomic structure
you are talking about. So if an atom has 8 protons (atomic number = 8), it must be oxygen. If an atom has 12 protons (atomic number = 12), it must be magnesium. Similarly, every chlorine atom (atomic number = 17) has 17 protons; every uranium atom (atomic number = 92) has 92 protons.
Isotopes The number of neutrons in an atom can vary within small limits. For example, there are three kinds of carbon atom 12C, 13C and 14C. They all have the same number of protons, but the number of neutrons varies.
carbon-12 carbon-13 carbon-14
protons 6 6 6
neutrons 6 7 8
mass number 12 13 14
These different atoms of carbon are called isotopes. The fact that they have varying numbers of neutrons makes no difference whatsoever to the chemical reactions of the carbon. Isotopes are atoms which have the same atomic number but different mass numbers. They have the same number of protons but different numbers of neutrons.
The electrons Working out the number of electrons Atoms are electrically neutral, and the positiveness of the protons is balanced by the negativeness of the electrons. It follows that in a neutral atom: no of electrons = no of protons
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a simple view of atomic structure
So, if an oxygen atom (atomic number = 8) has 8 protons, it must also have 8 electrons; if a chlorine atom (atomic number = 17) has 17 protons, it must also have 17 electrons. The arrangement of the electrons The electrons are found at considerable distances from the nucleus in a series of levels called energy levels. Each energy level can only hold a certain number of electrons. The first level (nearest the nucleus) will only hold 2 electrons, the second holds 8, and the third also seems to be full when it has 8 electrons. At GCSE you stop there because the pattern gets more complicated after that. These levels can be thought of as getting progressively further from the nucleus. Electrons will always go into the lowest possible energy level (nearest the nucleus) - provided there is space. To work out the electronic arrangement of an atom ●
●
●
Look up the atomic number in the Periodic Table - making sure that you choose the right number if two numbers are given. The atomic number will always be the smaller one. This tells you the number of protons, and hence the number of electrons. Arrange the electrons in levels, always filling up an inner level before you go to an outer one.
e.g. to find the electronic arrangement in chlorine ●
The Periodic Table gives you the atomic number of 17.
●
Therefore there are 17 protons and 17 electrons.
●
The arrangement of the electrons will be 2, 8, 7 (i.e. 2 in the first level, 8 in the second, and 7 in the third).
The electronic arrangements of the first 20 elements
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a simple view of atomic structure
After this the pattern alters as you enter the transition series in the Periodic Table. Two important generalisations If you look at the patterns in this table: ●
The number of electrons in the outer level is the same as the group number. (Except with helium which has only 2 electrons. The noble gases are also usually called group 0 - not group 8.) This pattern extends throughout the Periodic Table for the main groups (i.e. not including the transition elements). So if you know that barium is in group 2, it has 2 electrons in its outer level; iodine (group 7) has 7 electrons in its outer level; lead (group 4) has 4 electrons in its outer level.
●
Noble gases have full outer levels. This generalisation will need modifying for A'level purposes.
Dots-and-crosses diagrams In any introductory chemistry course you will have come across the electronic structures of hydrogen and carbon, for example, drawn as:
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a simple view of atomic structure
Note: There are many places where you could still make use of this model of the atom at A'level. It is, however, a simplification and can be misleading. It gives the impression that the electrons are circling the nucleus in orbits like planets around the sun. As you will find when you look at the A'level view of the atom, it is impossible to know exactly how they are actually moving.
The circles show energy levels - representing increasing distances from the nucleus. You could straighten the circles out and draw the electronic structure as a simple energy diagram. Carbon, for example, would look like this:
Thinking of the arrangement of the electrons in this way makes a useful bridge to the A'level view.
Note: If you have come to this page as a UK GCSE student (or a student on a similar introductory chemistry course elsewhere) and want some more help, you may be interested in my GCSE Chemistry book. This link will take you to a page describing it.
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a simple view of atomic structure
Where would you like to go now? To the atomic properties menu . . . To the atomic structure and bonding menu . . . To Main Menu . . .
© Jim Clark 2000
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GCSE chemistry
Understanding Chemistry
GCSE CHEMISTRY This book covers the chemistry content of all the UK GCSE Chemistry syllabuses - whether as a part of dual award science, or as a separate science. It is aimed at students likely to achieve grades from A* to B. If you are working in another system, GCSE in the UK is an exam taken at the end of a (usually) two year course at the age of 16. Anyone taking a similar introductory chemistry course may find the book helpful. On this page you will find a description of how the book is organised, together with a summary of the contents. You will also find direct links to the book on both the Longman and the Amazon.co.uk sites. Education in Chemistry, May 2003 "I was impressed with this new book, . . ." "The text is clearly laid out with excellent diagrams and illustrations." "This is an excellent textbook." School Science Review (issue 307) ". . will stretch and enthuse those with some ability in chemistry." http://www.chemguide.co.uk/gcsebook.html (1 of 5)30/12/2004 11:00:11
GCSE chemistry
"It would certainly help to bridge the gap between GCSE and AS level."
How to get hold of the book Schools or colleges would probably find it best to go to the Longman GCSE Chemistry website, but this site isn't really set up for individual purchases. You can, of course, buy the book through normal book sellers, but if you want to buy online, you will find a direct link to Amazon.co.uk coming up. Non-UK students can also buy the book from Amazon.co.uk, but will obviously have to pay a slightly higher delivery charge.
Note: If your usual source of books is Amazon.com, you should compare the price for the book (including delivery) from Amazon.com with the price from Amazon.co.uk - even if you live in North America. You may well find that it is significantly cheaper to buy from Amazon.co.uk and have it sent by air mail across the Atlantic, than to buy it in America!
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GCSE chemistry
Have a look at the book on the Amazon site
What the book covers The book is organised into 6 sections plus an important appendix. Each section is made up of a number of related chapters. There are questions at the end of each chapter to test understanding, and a set of GCSE-style exam questions at the end of each section. Answers to all the questions are provided on the supporting website although these are password-protected so that only teachers can get at them! Section A: Particles This covers an introduction to atomic structure and bonding (including metallic bonding and intermolecular forces) and the relationship between the structures of elements and compounds and their physical properties. There is also a chapter on how to write formulae and equations, and a final one on the factors affecting rates of reaction together with explanations. Section B: Some essential background chemistry This is a lengthy section which covers the important lab-based chemistry: ●
Reactivity series
●
Acids and their reactions
●
Salts
●
Simple analysis
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Periodic Table: including some history, the structure of the table, the noble gases, Groups 1 and 7, and an introduction to transition
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GCSE chemistry
metals ●
Electrolysis and electrochemical cells
●
Energy changes in reactions
Section C: Large scale chemistry This covers the extraction of several metals, and the chemistry of salt and limestone. It introduces reversible reactions leading to the Haber and Contact Processes. Section D: Air, water and earth Discusses the atmosphere (including its evolution and some environmental problems), water (including hardness, water treatment, and an introduction to colloids) and types of rock. Section E: Organic chemistry An introductory look at the oil industry and some simple organic compounds (alkanes, alkenes, alcohols, carboxylic acids, and a brief look at esters). Structural isomerism is explained where it arises. There are also chapters on food and drugs, and enzymes. Section F: Sums This section deals with all the calculations involving relative atomic masses and moles up to and including simple titration and electrolysis calculations. Appendices The most important appendix explains how to maximise your score when writing up coursework practical investigations to satisfy the requirements of UK GCSE examiners. The fully written out investigation is available from the website accompanying the book. (See below.) The website
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GCSE chemistry
There is a website to accompany the book which you can find by following this link. You may find this useful even if you don't end up buying the book! You will find lots of links to other other useful chemistry web sites, a fully written up example of a coursework investigation, and a set of worksheets. Answers to all the questions in the book are available, but only to teachers who have purchased the book from Longman. The answers are password-protected for obvious reasons!
Go to Main Menu . . .
© Jim Clark 2003
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atomic orbitals
ATOMIC ORBITALS
This page explains what an atomic orbital is. It explores s and p orbitals in some detail, including their shapes and energies. d orbitals are described only in terms of their energy, and f orbitals only get a passing mention.
What is an atomic orbital? Orbitals and orbits When the a planet moves around the sun, you can plot a definite path for it which is called an orbit. A simple view of the atom looks similar and you may have pictured the electrons as orbiting around the nucleus. The truth is different, and electrons in fact inhabit regions of space known as orbitals. Orbits and orbitals sound similar, but they have quite different meanings. It is essential that you understand the difference between them. The impossibility of drawing orbits for electrons To plot a path for something you need to know exactly where the object is and be able to work out exactly where it's going to be an instant later. You can't do this for electrons. The Heisenberg Uncertainty Principle (not required at A'level) says loosely - that you can't know with certainty both where an electron is and where it's going next. That makes it impossible to plot an orbit for an electron around a nucleus. Is this a big problem? No. If something is impossible, you have to accept it and find a way around it. Hydrogen's electron - the 1s orbital
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atomic orbitals
Note: In this diagram (and the orbital diagrams that follow), the nucleus is shown very much larger than it really is. This is just for clarity.
Suppose you had a single hydrogen atom and at a particular instant plotted the position of the one electron. Soon afterwards, you do the same thing, and find that it is in a new position. You have no idea how it got from the first place to the second. You keep on doing this over and over again, and gradually build up a sort of 3D map of the places that the electron is likely to be found. In the hydrogen case, the electron can be found anywhere within a spherical space surrounding the nucleus. The diagram shows a crosssection through this spherical space. 95% of the time (or any other percentage you choose), the electron will be found within a fairly easily defined region of space quite close to the nucleus. Such a region of space is called an orbital. You can think of an orbital as being the region of space in which the electron lives.
Note: If you wanted to be absolutely 100% sure of where the electron is, you would have to draw an orbital the size of the Universe!
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atomic orbitals
What is the electron doing in the orbital? We don't know, we can't know, and so we just ignore the problem! All you can say is that if an electron is in a particular orbital it will have a particular definable energy. Each orbital has a name. The orbital occupied by the hydrogen electron is called a 1s orbital. The "1" represents the fact that the orbital is in the energy level closest to the nucleus. The "s" tells you about the shape of the orbital. s orbitals are spherically symmetric around the nucleus - in each case, like a hollow ball made of rather chunky material with the nucleus at its centre. The orbital on the left is a 2s orbital. This is similar to a 1s orbital except that the region where there is the greatest chance of finding the electron is further from the nucleus - this is an orbital at the second energy level. If you look carefully, you will notice that there is another region of slightly higher electron density (where the dots are thicker) nearer the nucleus. ("Electron density" is another way of talking about how likely you are to find an electron at a particular place.) 2s (and 3s, 4s, etc) electrons spend some of their time closer to the nucleus than you might expect. The effect of this is to slightly reduce the energy of electrons in s orbitals. The nearer the nucleus the electrons get, the lower their energy. 3s, 4s (etc) orbitals get progressively further from the nucleus. p orbitals
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atomic orbitals
Not all electrons inhabit s orbitals (in fact, very few electrons live in s orbitals). At the first energy level, the only orbital available to electrons is the 1s orbital, but at the second level, as well as a 2s orbital, there are also orbitals called 2p orbitals. A p orbital is rather like 2 identical balloons tied together at the nucleus. The diagram on the right is a cross-section through that 3-dimensional region of space. Once again, the orbital shows where there is a 95% chance of finding a particular electron.
Beyond A'level: If you imagine a horizontal plane through the nucleus, with one lobe of the orbital above the plane and the other beneath it, there is a zero probability of finding the electron on that plane. So how does the electron get from one lobe to the other if it can never pass through the plane of the nucleus? For A'level chemistry you just have to accept that it does! If you want to find out more, read about the wave nature of electrons.
Unlike an s orbital, a p orbital points in a particular direction - the one drawn points up and down the page. At any one energy level it is possible to have three absolutely equivalent p orbitals pointing mutually at right angles to each other. These are arbitrarily given the symbols px, py and pz. This is simply for convenience - what you might think of as the x, y or z direction changes constantly as the atom tumbles in space.
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atomic orbitals
The p orbitals at the second energy level are called 2px, 2py and 2pz. There are similar orbitals at subsequent levels - 3px, 3py, 3pz, 4px, 4py, 4pz and so on. All levels except for the first level have p orbitals. At the higher levels the lobes get more elongated, with the most likely place to find the electron more distant from the nucleus.
d and f orbitals In addition to s and p orbitals, there are two other sets of orbitals which become available for electrons to inhabit at higher energy levels. At the third level, there is a set of five d orbitals (with complicated shapes and names) as well as the 3s and 3p orbitals (3px, 3py, 3pz). At the third level there are a total of nine orbitals altogether. At the fourth level, as well the 4s and 4p and 4d orbitals there are an additional seven f orbitals - 16 orbitals in all. s, p, d and f orbitals are then available at all higher energy levels as well. For A'level purposes, you have to be aware that there are sets of five d orbitals at levels from the third level upwards, but you will not be expected to draw them or name them. Apart from a passing reference, you won't come across f orbitals at all.
Fitting electrons into orbitals You can think of an atom as a very bizarre house (like an inverted pyramid!) - with the nucleus living on the ground floor, and then various rooms (orbitals) on the higher floors occupied by the electrons. On the first floor there is only 1 room (the 1s orbital); on the second floor there are 4 rooms (the 2s, 2px, 2py and 2pz orbitals); on the third floor there are 9 rooms (one 3s orbital, three 3p orbitals and five 3d orbitals); and so on. But the rooms aren't very big . . . Each orbital can only hold 2 http://www.chemguide.co.uk/atoms/properties/atomorbs.html (5 of 8)30/12/2004 11:00:17
atomic orbitals
electrons. A convenient way of showing the orbitals that the electrons live in is to draw "electrons-in-boxes". "Electrons-in-boxes" Orbitals can be represented as boxes with the electrons in them shown as arrows. Often an up-arrow and a down-arrow are used to show that the electrons are in some way different.
Beyond A'level: The need to have all electrons in an atom different comes out of quantum theory. If they live in different orbitals, that's fine - but if they are both in the same orbital there has to be some subtle distinction between them. Quantum theory allocates them a property known as "spin" which is what the arrows are intended to suggest.
A 1s orbital holding 2 electrons would be drawn as shown on the right, but it can be written even more quickly as 1s2. This is read as "one s two" - not as "one s squared". You mustn't confuse the two numbers in this notation:
The order of filling orbitals Electrons fill low energy orbitals (closer to the nucleus) before they fill higher energy ones. Where there is a choice between orbitals of equal energy, they fill the orbitals singly as far as possible. This filling of orbitals singly where possible is known as Hund's rule. It only applies where the orbitals have exactly the same energies (as with p orbitals, for example), and helps to minimise the repulsions between electrons and so makes the atom more stable.
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atomic orbitals
The diagram (not to scale) summarises the energies of the orbitals up to the 4p level.
Notice that the s orbital always has a slightly lower energy than the p orbitals at the same energy level, so the s orbital always fills with electrons before the corresponding p orbitals. The real oddity is the position of the 3d orbitals. They are at a slightly higher level than the 4s - and so it is the 4s orbital which will fill first, followed by all the 3d orbitals and then the 4p orbitals. Similar confusion occurs at higher levels, with so much overlap between the energy levels that the 4f orbitals don't fill until after the 6s, for example. For A'level purposes you simply have to remember that the 4s orbital fills before the 3d orbitals. The same thing happens at the next level as well the 5s orbital fills before the 4d orbitals. All the other complications are beyond A'level. Knowing the order of filling is central to understanding how to write electronic structures. Follow the link below to find out how to do this.
Where would you like to go now? To look at how to write electronic structures . . . To the atomic properties menu . . . http://www.chemguide.co.uk/atoms/properties/atomorbs.html (7 of 8)30/12/2004 11:00:17
atomic orbitals
To the atomic structure and bonding menu . . . To Main Menu . . .
© Jim Clark 2000 (modified 2004)
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electronic structures of atoms
ELECTRONIC STRUCTURES
This page explores how you write electronic structures for atoms using s, p, and d notation. It assumes that you know about simple atomic orbitals - at least as far as the way they are named, and their relative energies. If you want to look at the electronic structures of simple monatomic ions (such as Cl-, Ca2+ and Cr3+), you will find a link at the bottom of the page.
Important! If you haven't already read the page on atomic orbitals you should follow this link before you go any further.
The electronic structures of atoms Relating orbital filling to the Periodic Table
Most A'level syllabuses stop at krypton when it comes to writing electronic structures, but it is possible that you could be asked for structures for elements up as far as barium. After barium you have to worry about f orbitals as well as s, p and d orbitals - and that's a problem beyond A'level. It is important that you look through past exam papers as well as your syllabus so that you can judge how hard the questions are likely to get. This page looks in detail at the elements in the shortened version of the Periodic Table above, and then shows how you could work out the structures of some bigger atoms.
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electronic structures of atoms
Important! You must have a copy of your syllabus and copies of recent exam papers. If you haven't got them, follow this link to find out how to get hold of them.
The first period Hydrogen has its only electron in the 1s orbital - 1s1, and at helium the first level is completely full - 1s2. The second period Now we need to start filling the second level, and hence start the second period. Lithium's electron goes into the 2s orbital because that has a lower energy than the 2p orbitals. Lithium has an electronic structure of 1s22s1. Beryllium adds a second electron to this same level - 1s22s2. Now the 2p levels start to fill. These levels all have the same energy, and so the electrons go in singly at first. B
1s22s22px1
C
1s22s22px12py1
N
1s22s22px12py12pz1 Note: The orbitals where something new is happening are shown in bold type. You wouldn't normally write them any differently from the other orbitals.
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electronic structures of atoms
The next electrons to go in will have to pair up with those already there. O
1s22s22px22py12pz1
F
1s22s22px22py22pz1
Ne
1s22s22px22py22pz2
You can see that it is going to get progressively tedious to write the full electronic structures of atoms as the number of electrons increases. There are two ways around this, and you must be familiar with both. Shortcut 1: All the various p electrons can be lumped together. For example, fluorine could be written as 1s22s22p5, and neon as 1s22s22p6. This is what is normally done if the electrons are in an inner layer. If the electrons are in the bonding level (those on the outside of the atom), they are sometimes written in shorthand, sometimes in full. Don't worry about this. Be prepared to meet either version, but if you are asked for the electronic structure of something in an exam, write it out in full showing all the px, py and pz orbitals in the outer level separately. For example, although we haven't yet met the electronic structure of chlorine, you could write it as 1s22s22p63s23px23py23pz1. Notice that the 2p electrons are all lumped together whereas the 3p ones are shown in full. The logic is that the 3p electrons will be involved in bonding because they are on the outside of the atom, whereas the 2p electrons are buried deep in the atom and aren't really of any interest. Shortcut 2: You can lump all the inner electrons together using, for example, the symbol [Ne]. In this context, [Ne] means the electronic structure of neon - in other words: 1s22s22px22py22pz2 You wouldn't do this with helium because it takes longer to write [He] than it does 1s2. On this basis the structure of chlorine would be written [Ne] 3s23px23py23pz1. The third period
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electronic structures of atoms
At neon, all the second level orbitals are full, and so after this we have to start the third period with sodium. The pattern of filling is now exactly the same as in the previous period, except that everything is now happening at the 3-level. For example: short version Mg
1s22s22p63s2
[Ne]3s2
S
1s22s22p63s23px23py13pz1
[Ne]3s23px23py13pz1
Ar
1s22s22p63s23px23py23pz2
[Ne]3s23px23py23pz2
Note: Check that you can do these. Cover the text and then work out these structures for yourself. Then do all the rest of this period. When you've finished, check your answers against the corresponding elements from the previous period. Your answers should be the same except a level further out.
The beginning of the fourth period At this point the 3-level orbitals aren't all full - the 3d levels haven't been used yet. But if you refer back to the energies of the orbitals, you will see that the next lowest energy orbital is the 4s - so that fills next. K
1s22s22p63s23p64s1
Ca
1s22s22p63s23p64s2
There is strong evidence for this in the similarities in the chemistry of elements like sodium (1s22s22p63s1) and potassium (1s22s22p63s23p64s1) The outer electron governs their properties and that electron is in the same sort of orbital in both of the elements. That wouldn't be true if the outer electron in potassium was 3d1. s- and p-block elements
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electronic structures of atoms
The elements in group 1 of the Periodic Table all have an outer electronic structure of ns1 (where n is a number between 2 and 7). All group 2 elements have an outer electronic structure of ns2. Elements in groups 1 and 2 are described as s-block elements. Elements from group 3 across to the noble gases all have their outer electrons in p orbitals. These are then described as p-block elements. d-block elements
Remember that the 4s orbital has a lower energy than the 3d orbitals and so fills first. Once the 3d orbitals have filled up, the next electrons go into the 4p orbitals as you would expect. d-block elements are elements in which the last electron to be added to the atom is in a d orbital. The first series of these contains the elements from scandium to zinc, which at GCSE you probably called transition elements or transition metals. The terms "transition element" and "d-
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electronic structures of atoms
block element" don't quite have the same meaning, but it doesn't matter in the present context.
If you are interested: A transition element is defined as one which has partially filled d orbitals either in the element or any of its compounds. Zinc (at the right-hand end of the dblock) always has a completely full 3d level (3d10) and so doesn't count as a transition element.
d electrons are almost always described as, for example, d5 or d8 - and not written as separate orbitals. Remember that there are five d orbitals, and that the electrons will inhabit them singly as far as possible. Up to 5 electrons will occupy orbitals on their own. After that they will have to pair up. d5 means
d8 means
Notice in what follows that all the 3-level orbitals are written together, even though the 3d electrons are added to the atom after the 4s. Sc
1s22s22p63s23p63d14s2
Ti
1s22s22p63s23p63d24s2
V
1s22s22p63s23p63d34s2
Cr
1s22s22p63s23p63d54s1
Whoops! Chromium breaks the sequence. In chromium, the electrons in the 3d and 4s orbitals rearrange so that there is one electron in each orbital. It would be convenient if the sequence was tidy - but it's not! Mn
1s22s22p63s23p63d54s2
Fe
1s22s22p63s23p63d64s2
(back to being tidy again)
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electronic structures of atoms
Co
1s22s22p63s23p63d74s2
Ni
1s22s22p63s23p63d84s2
Cu
1s22s22p63s23p63d104s1
Zn
1s22s22p63s23p63d104s2
(another awkward one!)
And at zinc the process of filling the d orbitals is complete. Filling the rest of period 4 The next orbitals to be used are the 4p, and these fill in exactly the same way as the 2p or 3p. We are back now with the p-block elements from gallium to krypton. Bromine, for example, is 1s22s22p63s23p63d104s24px24py24pz1.
Useful exercise: Work out the electronic structures of all the elements from gallium to krypton. You can check your answers by comparing them with the elements directly above them in the Periodic Table. For example, gallium will have the same sort of arrangement of its outer level electrons as boron or aluminium - except that gallium's outer electrons will be in the 4-level.
Summary Writing the electronic structure of an element from hydrogen to krypton ●
●
●
Use the Periodic Table to find the atomic number, and hence number of electrons. Fill up orbitals in the order 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p - until you run out of electrons. The 3d is the awkward one - remember that specially. Fill p and d orbitals singly as far as possible before pairing electrons up. Remember that chromium and copper have electronic structures which break the pattern in the first row of the d-block.
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Writing the electronic structure of big s- or p-block elements
Note: We are deliberately excluding the d-block elements apart from the first row that we've already looked at in detail. The pattern of awkward structures isn't the same in the other rows. This isn't an A'level problem.
First work out the number of outer electrons. This is quite likely all you will be asked to do anyway. The number of outer electrons is the same as the group number. (The noble gases are a bit of a problem here, because they are normally called group 0 rather then group 8. Helium has 2 outer electrons; the rest have 8.) All elements in group 3, for example, have 3 electrons in their outer level. Fit these electrons into s and p orbitals as necessary. Which level orbitals? Count the periods in the Periodic Table (not forgetting the one with H and He in it). Iodine is in group 7 and so has 7 outer electrons. It is in the fifth period and so its electrons will be in 5s and 5p orbitals. Iodine has the outer structure 5s25px25py25pz1. What about the inner electrons if you need to work them out as well? The 1, 2 and 3 levels will all be full, and so will the 4s, 4p and 4d. The 4f levels don't fill until after anything you will be asked about at A'level. Just forget about them! That gives the full structure: 1s22s22p63s23p63d104s24p64d105s25px25py25pz1. When you've finished, count all the electrons to make sure that they come to the same as the atomic number. Don't forget to make this check - it's easy to miss an orbital out when it gets this complicated. Barium is in group 2 and so has 2 outer electrons. It is in the sixth period. Barium has the outer structure 6s2. Including all the inner levels: 1s22s22p63s23p63d104s24p64d105s25p66s2. http://www.chemguide.co.uk/atoms/properties/elstructs.html (8 of 9)30/12/2004 11:00:22
electronic structures of atoms
It would be easy to include 5d10 as well by mistake, but the d level always fills after the next s level - so 5d fills after 6s just as 3d fills after 4s. As long as you counted the number of electrons you could easily spot this mistake because you would have 10 too many.
Note: Don't worry too much about these complicated structures. You need to know how to work them out in principle, but your examiners are much more likely to ask you for something simple like sulphur or iron.
Where would you like to go now? To working out electronic structures for ions . . . To the atomic properties menu . . . To the atomic structure and bonding menu . . . To Main Menu . . .
© Jim Clark 2000
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Understanding Chemistry - A' level syllabuses
Understanding Chemistry
UK A and AS LEVEL CHEMISTRY SYLLABUSES I assume that you want to get the best grade you possibly can with the minimum of effort! Getting a good A level grade is rather like playing a game with your examiners - in which they make up the rules (and occasionally change them). You aren't going to win unless you know those rules. Before you do anything else: ●
●
Get a copy of your syllabus if you haven't already got one. Details of how to do this are given below. Syllabuses are often quite difficult to interpret, so you need to know exactly what questions your examiners are asking, and how they are marking them. Explore your Exam Board web site. They all offer free downloads of specimen papers (including mark schemes), but you might have to pay for recent exam papers and mark schemes, and other support material. If they don't offer these free, find out how to order them and invest a small amount of money in your future! If you want the best possible grade, you should be working with exam papers all the way through your course. Leaving looking at exam papers until your last minute revision is too late. Be careful, though! Syllabuses change and so do examiners. Make sure that the question papers and mark schemes you get relate to your current syllabus and are as recent as possible. A new chief examiner can make a lot of difference to the style of a question paper.
How to download a copy of your syllabus
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Understanding Chemistry - A' level syllabuses
Finding your way to the right syllabus The following links take you to the front pages of each of the Exam Board web sites and you will then have to find your own way to your syllabus. This is because these sites are liable to change. Be aware that the syllabuses are known as specifications. You want GCE Advanced and Advanced Subsidiary (A and AS) Chemistry. Finding the syllabuses is very straightforward - finding other information may take you longer! The Exam Boards: ●
OCR This includes both the standard OCR syllabus and the Salters syllabus.
●
Edexcel This includes both the standard Edexcel syllabus and the Nuffield syllabus.
●
AQA AQA have free downloadable versions of all their recent exam papers and mark schemes. Once you get to the chemistry page, look for it under "Assessment Material". You can also get Examiners' Reports (another link from the chemistry page). These are essential if you want to avoid common mistakes.
●
WJEC This link should take you directly to the correct chemistry page to download a syllabus. At the time of writing, you will have to pay if you want past papers or mark schemes.
Problems reading the downloaded syllabus? The syllabuses are available only in pdf format. You need software such as Adobe Acrobat Reader to access it. You have almost certainly got this http://www.chemguide.co.uk/syllabuses.html (2 of 3)30/12/2004 11:00:23
Understanding Chemistry - A' level syllabuses
(or the equivalent) software on your computer, but if your computer is old, you may not have the latest version. If your downloaded syllabus won't open, it may be that the syllabus was created in a newer version of the Reader than you've got. You will have to download a new version of Reader. Each of the Exam Board web sites provides a link to Adobe, but these links are often easy to miss. Use this link: ●
www.adobe.com This will take you to Adobe's front page where you will find a link enabling you to download the Reader. Be warned that this is a seriously large bit of software and could take a long time to download on a dial-up connection.
Go to Main Menu . . .
© Jim Clark 2000 (modified 2004)
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electronic structures of ions
ELECTRONIC STRUCTURES OF IONS
This page explores how you write electronic structures for simple monatomic ions (ions containing only one atom) using s, p, and d notation. It assumes that you already understand how to write electronic structures for atoms.
Important! If you have come straight to this page via a search engine, you should read the page on electronic structures of atoms before you go any further.
Working out the electronic structures of ions Ions are atoms (or groups of atoms) which carry an electric charge because they have either gained or lost one or more electrons. If an atom gains electrons it acquires a negative charge. If it loses electrons, it becomes positively charged. The electronic structure of s- and p-block ions Write the electronic structure for the neutral atom, and then add (for a negative ion) or subtract electrons (for a positive ion). To write the electronic structure for Cl -: Cl
1s22s22p63s23px23py23pz1
Cl-
1s22s22p63s23px23py23pz2
but Cl- has one more electron
To write the electronic structure for O2-: O
1s22s22px22py12pz1
but O2- has two more electrons
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O2-
1s22s22px22py22pz2
To write the electronic structure for Na+: Na
1s22s22p63s1
Na+
1s22s22p6
but Na+ has one less electron
To write the electronic structure for Ca2+: Ca
1s22s22p63s23p64s2
Ca2+
1s22s22p63s23p6
but Ca2+ has two less electrons
The electronic structure of d-block ions Here you are faced with one of the most irritating facts in A'level chemistry! You will recall that the first transition series (from scandium to zinc) is the result of the 3d orbitals being filled after the 4s orbital. However, once the electrons are established in their orbitals, the energy order changes - and in all the chemistry of the transition elements, the 4s orbital behaves as the outermost, highest energy orbital. The reversed order of the 3d and 4s orbitals only applies to building the atom up in the first place. In all other respects, the 4s electrons are always the electrons you need to think about first. You must remember this:
When d-block elements form ions, the 4s electrons are lost first.
Provided you remember that, working out the structure of a d-block ion is no different from working out the structure of, say, a sodium ion.
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electronic structures of ions
To write the electronic structure for Cr3+: Cr
1s22s22p63s23p63d54s1
Cr3+
1s22s22p63s23p63d3
The 4s electron is lost first followed by two of the 3d electrons.
To write the electronic structure for Zn2+: Zn
1s22s22p63s23p63d104s2
Zn2+
1s22s22p63s23p63d10
This time there is no need to use any of the 3d electrons.
To write the electronic structure for Fe3+: Fe
1s22s22p63s23p63d64s2
Fe3+
1s22s22p63s23p63d5
The 4s electrons are lost first followed by one of the 3d electrons. The rule is quite simple. Take the 4s electrons off first, and then as many 3d electrons as necessary to produce the correct positive charge.
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electronic structures of ions
Note: You may well have the impression from GCSE that ions have to have noble gas structures. It's not true! Most (but not all) ions formed by s- and p-block elements do have noble gas structures, but if you look at the d-block ions we've used as examples, not one of them has a noble gas structure - yet they are all perfectly valid ions. Getting away from a reliance on the concept of noble gas structures is one of the difficult mental leaps that you have to make at the beginning of A'level chemistry.
Where would you like to go now? To the atomic properties menu . . . To the atomic structure and bonding menu . . . To Main Menu . . .
© Jim Clark 2000
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first ionisation energy
IONISATION ENERGY
This page explains what first ionisation energy is, and then looks at the way it varies around the Periodic Table - across periods and down groups. It assumes that you know about simple atomic orbitals, and can write electronic structures for simple atoms. You will find a link at the bottom of the page to a similar description of successive ionisation energies (second, third and so on).
Important! If you aren't reasonable happy about atomic orbitals and electronic structures you should follow these links before you go any further.
Defining first ionisation energy Definition The first ionisation energy is the energy required to remove the most loosely held electron from one mole of gaseous atoms to produce 1 mole of gaseous ions each with a charge of 1+. This is more easily seen in symbol terms.
It is the energy needed to carry out this change per mole of X.
Worried about moles? Don't be! For now, just take it as a measure of a particular amount of a substance. It isn't worth worrying about at the moment.
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first ionisation energy
Things to notice about the equation The state symbols - (g) - are essential. When you are talking about ionisation energies, everything must be present in the gas state. Ionisation energies are measured in kJ mol-1 (kilojoules per mole). They vary in size from 381 (which you would consider very low) up to 2370 (which is very high). All elements have a first ionisation energy - even atoms which don't form positive ions in test tubes. The reason that helium (1st I.E. = 2370 kJ mol1) doesn't normally form a positive ion is because of the huge amount of energy that would be needed to remove one of its electrons.
Patterns of first ionisation energies in the Periodic Table The first 20 elements
First ionisation energy shows periodicity. That means that it varies in a repetitive way as you move through the Periodic Table. For example, look at the pattern from Li to Ne, and then compare it with the identical pattern from Na to Ar. These variations in first ionisation energy can all be explained in terms of the structures of the atoms involved.
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first ionisation energy
Ionisation energy is a measure of the energy needed to pull a particular electron away from the attraction of the nucleus. A high value of ionisation energy shows a high attraction between the electron and the nucleus. The size of that attraction will be governed by: The charge on the nucleus. The more protons there are in the nucleus, the more positively charged the nucleus is, and the more strongly electrons are attracted to it. The distance of the electron from the nucleus. Attraction falls off very rapidly with distance. An electron close to the nucleus will be much more strongly attracted than one further away. The number of electrons between the outer electrons and the nucleus. Consider a sodium atom, with the electronic structure 2,8,1. (There's no reason why you can't use this notation if it's useful!) If the outer electron looks in towards the nucleus, it doesn't see the nucleus sharply. Between it and the nucleus there are the two layers of electrons in the first and second levels. The 11 protons in the sodium's nucleus have their effect cut down by the 10 inner electrons. The outer electron therefore only feels a net pull of approximately 1+ from the centre. This lessening of the pull of the nucleus by inner electrons is known as screening or shielding.
Warning! Electrons don't, of course, "look in" towards the nucleus - and they don't "see" anything either! But there's no reason why you can't imagine it in these terms if it helps you to visualise what's happening. Just don't use these terms in an exam! You may get an examiner who is upset by this sort of loose language.
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Whether the electron is on its own in an orbital or paired with another electron. Two electrons in the same orbital experience a bit of repulsion from each other. This offsets the attraction of the nucleus, so that paired electrons are removed rather more easily than you might expect.
Explaining the pattern in the first few elements Hydrogen has an electronic structure of 1s1. It is a very small atom, and the single electron is close to the nucleus and therefore strongly attracted. There are no electrons screening it from the nucleus and so the ionisation energy is high (1310 kJ mol-1). Helium has a structure 1s2. The electron is being removed from the same orbital as in hydrogen's case. It is close to the nucleus and unscreened. The value of the ionisation energy (2370 kJ mol-1) is much higher than hydrogen, because the nucleus now has 2 protons attracting the electrons instead of 1. Lithium is 1s22s1. Its outer electron is in the second energy level, much more distant from the nucleus. You might argue that that would be offset by the additional proton in the nucleus, but the electron doesn't feel the full pull of the nucleus - it is screened by the 1s2 electrons.
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You can think of the electron as feeling a net 1+ pull from the centre (3 protons offset by the two 1s2 electrons). If you compare lithium with hydrogen (instead of with helium), the hydrogen's electron also feels a 1+ pull from the nucleus, but the distance is much greater with lithium. Lithium's first ionisation energy drops to 519 kJ mol-1 whereas hydrogen's is 1310 kJ mol-1.
The patterns in periods 2 and 3 Talking through the next 17 atoms one at a time would take ages. We can do it much more neatly by explaining the main trends in these periods, and then accounting for the exceptions to these trends. The first thing to realise is that the patterns in the two periods are identical - the difference being that the ionisation energies in period 3 are all lower than those in period 2.
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Explaining the general trend across periods 2 and 3 The general trend is for ionisation energies to increase across a period. In the whole of period 2, the outer electrons are in 2-level orbitals - 2s or 2p. These are all the same sort of distances from the nucleus, and are screened by the same 1s2 electrons. The major difference is the increasing number of protons in the nucleus as you go from lithium to neon. That causes greater attraction between the nucleus and the electrons and so increases the ionisation energies. In fact the increasing nuclear charge also drags the outer electrons in closer to the nucleus. That increases ionisation energies still more as you go across the period.
Note: Factors affecting atomic radius are covered on a separate page.
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In period 3, the trend is exactly the same. This time, all the electrons being removed are in the third level and are screened by the 1s22s22p6 electrons. They all have the same sort of environment, but there is an increasing nuclear charge. Why the drop between groups 2 and 3 (Be-B and Mg-Al)? The explanation lies with the structures of boron and aluminium. The outer electron is removed more easily from these atoms than the general trend in their period would suggest. Be
1s22s2
1st I.E. = 900 kJ mol-1
B
1s22s22px1
1st I.E. = 799 kJ mol-1
You might expect the boron value to be more than the beryllium value because of the extra proton. Offsetting that is the fact that boron's outer electron is in a 2p orbital rather than a 2s. 2p orbitals have a slightly higher energy than the 2s orbital, and the electron is, on average, to be found further from the nucleus. This has two effects. ●
●
The increased distance results in a reduced attraction and so a reduced ionisation energy. The 2p orbital is screened not only by the 1s2 electrons but, to some extent, by the 2s2 electrons as well. That also reduces the pull from the nucleus and so lowers the ionisation energy.
The explanation for the drop between magnesium and aluminium is the same, except that everything is happening at the 3-level rather than the 2level. Mg
1s22s22p63s2
1st I.E. = 736 kJ mol-1
Al
1s22s22p63s23px1
1st I.E. = 577 kJ mol-1
The 3p electron in aluminium is slightly more distant from the nucleus than the 3s, and partially screened by the 3s2 electrons as well as the inner electrons. Both of these factors offset the effect of the extra proton.
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first ionisation energy
Warning! You might possibly come across a text book which describes the drop between group 2 and group 3 by saying that a full s2 orbital is in some way especially stable and that makes the electron more difficult to remove. In other words, that the fluctuation is because the group 2 value for ionisation energy is abnormally high. This is quite simply wrong! The reason for the fluctuation is because the group 3 value is lower than you might expect for the reasons we've looked at.
Why the drop between groups 5 and 6 (N-O and P-S)? Once again, you might expect the ionisation energy of the group 6 element to be higher than that of group 5 because of the extra proton. What is offsetting it this time? N
1s22s22px12py12pz1
1st I.E. = 1400 kJ mol-1
O
1s22s22px22py12pz1
1st I.E. = 1310 kJ mol-1
The screening is identical (from the 1s2 and, to some extent, from the 2s2 electrons), and the electron is being removed from an identical orbital. The difference is that in the oxygen case the electron being removed is one of the 2px2 pair. The repulsion between the two electrons in the same orbital means that the electron is easier to remove than it would otherwise be. The drop in ionisation energy at sulphur is accounted for in the same way.
Trends in ionisation energy down a group As you go down a group in the Periodic Table ionisation energies generally fall. You have already seen evidence of this in the fact that the ionisation energies in period 3 are all less than those in period 2. Taking Group 1 as a typical example:
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first ionisation energy
Why is the sodium value less than that of lithium? There are 11 protons in a sodium atom but only 3 in a lithium atom, so the nuclear charge is much greater. You might have expected a much larger ionisation energy in sodium, but offsetting the nuclear charge is a greater distance from the nucleus and more screening. Li
1s22s1
1st I.E. = 519 kJ mol-1
Na
1s22s22p63s1
1st I.E. = 494 kJ mol-1
Lithium's outer electron is in the second level, and only has the 1s2 electrons to screen it. The 2s1 electron feels the pull of 3 protons screened by 2 electrons - a net pull from the centre of 1+. The sodium's outer electron is in the third level, and is screened from the 11 protons in the nucleus by a total of 10 inner electrons. The 3s1 electron also feels a net pull of 1+ from the centre of the atom. In other words, the effect of the extra protons is compensated for by the effect of the extra screening electrons. The only factor left is the extra distance between the outer electron and the nucleus in sodium's case. That lowers the ionisation energy. Similar explanations hold as you go down the rest of this group - or, indeed, any other group.
Trends in ionisation energy in a transition series
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Apart from zinc at the end, the other ionisation energies are all much the same. All of these elements have an electronic structure [Ar]3dn4s2 (or 4s1 in the cases of chromium and copper). The electron being lost always comes from the 4s orbital.
Note: Confusingly, once the orbitals have electrons in them, the 4s orbital has a higher energy than the 3d - quite the opposite of their order when the atoms are being filled with electrons. That means that it is a 4s electron which is lost from the atom when it forms an ion. It also means that the 3d orbitals are slightly closer to the nucleus than the 4s - and so offer some screening. You will find this commented on in the page about electronic structures of ions.
As you go from one atom to the next in the series, the number of protons in the nucleus increases, but so also does the number of 3d electrons. The 3d electrons have some screening effect, and the extra proton and the extra 3d electron more or less cancel each other out as far as attraction from the centre of the atom is concerned. The rise at zinc is easy to explain. Cu
[Ar]3d104s1
1st I.E. = 745 kJ mol-1
Zn
[Ar]3d104s2
1st I.E. = 908 kJ mol-1
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first ionisation energy
screening, but the zinc has one extra proton in the nucleus and so the attraction is greater.
Ionisation energies and reactivity The lower the ionisation energy, the more easily this change happens:
You can explain the increase in reactivity of the Group 1 metals (Li, Na, K, Rb, Cs) as you go down the group in terms of the fall in ionisation energy. Whatever these metals react with, they have to form positive ions in the process, and so the lower the ionisation energy, the more easily those ions will form. The danger with this approach is that the formation of the positive ion is only one stage in a multi-step process. For example, you wouldn't be starting with gaseous atoms; nor would you end up with gaseous positive ions - you would end up with ions in a solid or in solution. The energy changes in these processes also vary from element to element. Ideally you need to consider the whole picture and not just one small part of it. However, the ionisation energies of the elements are going to be major contributing factors towards the activation energy of the reactions. Remember that activation energy is the minimum energy needed before a reaction will take place. The lower the activation energy, the faster the reaction will be - irrespective of what the overall energy changes in the reaction are. The fall in ionisation energy as you go down a group will lead to lower activation energies and therefore faster reactions.
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Note: You will find a page discussing this in more detail in the inorganic section of this site dealing with the reactions of Group 2 metals with water.
Where would you like to go now? To look at second (and successive) ionisation energies . . . To the atomic properties menu . . . To the atomic structure and bonding menu . . . To Main Menu . . .
© Jim Clark 2000
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atomic and ionic radius
ATOMIC AND IONIC RADIUS
This page explains the various measures of atomic radius, and then looks at the way it varies around the Periodic Table - across periods and down groups. It assumes that you understand electronic structures for simple atoms written in s, p, d notation.
Important! If you aren't reasonable happy about electronic structures you should follow this link before you go any further.
ATOMIC RADIUS Measures of atomic radius Unlike a ball, an atom doesn't have a fixed radius. The radius of an atom can only be found by measuring the distance between the nuclei of two touching atoms, and then halving that distance.
As you can see from the diagrams, the same atom could be found to have a different radius depending on what was around it. The left hand diagram shows bonded atoms. The atoms are pulled closely together and so the measured radius is less than if they are just touching. This is what you would get if you had metal atoms in a metallic structure, or atoms covalently bonded to each other. The type of atomic radius being measured here is called the metallic radius or the covalent radius depending on the bonding.
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atomic and ionic radius
The right hand diagram shows what happens if the atoms are just touching. The attractive forces are much less, and the atoms are essentially "unsquashed". This measure of atomic radius is called the van der Waals radius after the weak attractions present in this situation.
Note: If you want to explore these various types of bonding this link will take you to the bonding menu.
Trends in atomic radius in the Periodic Table The exact pattern you get depends on which measure of atomic radius you use - but the trends are still valid. The following diagram uses metallic radii for metallic elements, covalent radii for elements that form covalent bonds, and van der Waals radii for those (like the noble gases) which don't form bonds. Trends in atomic radius in Periods 2 and 3
Trends in atomic radius down a group It is fairly obvious that the atoms get bigger as you go down groups. The reason is equally obvious - you are adding extra layers of electrons. Trends in atomic radius across periods You have to ignore the noble gas at the end of each period. Because neon and argon don't form bonds, you can only measure their van der Waals radius - a case where the atom is pretty well "unsquashed". All the other atoms are being measured where their atomic radius is being lessened by strong attractions. You aren't comparing like with like if you
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include the noble gases.
Leaving the noble gases out, atoms get smaller as you go across a period.
If you think about it, the metallic or covalent radius is going to be a measure of the distance from the nucleus to the electrons which make up the bond. (Look back to the left-hand side of the first diagram on this page if you aren't sure, and picture the bonding electrons as being half way between the two nuclei.) From lithium to fluorine, those electrons are all in the 2-level, being screened by the 1s2 electrons. The increasing number of protons in the nucleus as you go across the period pulls the electrons in more tightly. The amount of screening is constant for all of these elements.
Note: You might possibly wonder why you don't get extra screening from the 2s2 electrons in the cases of the elements from boron to fluorine where the bonding involves the p electrons. In each of these cases, before bonding happens, the existing s and p orbitals are reorganised (hybridised) into new orbitals of equal energy. When these atoms are bonded, there aren't any 2s electrons as such. If you don't know about hybridisation, just ignore this comment - you won't need it for UK A level purposes anyway.
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In the period from sodium to chlorine, the same thing happens. The size of the atom is controlled by the 3-level bonding electrons being pulled closer to the nucleus by increasing numbers of protons - in each case, screened by the 1- and 2-level electrons.
Trends in the transition elements
Although there is a slight contraction at the beginning of the series, the atoms are all much the same size. The size is determined by the 4s electrons. The pull of the increasing number of protons in the nucleus is more or less offset by the extra screening due to the increasing number of 3d electrons.
Note: Confusingly, once the orbitals have electrons in them, the 4s orbital has a higher energy than the 3d - quite the opposite of their order when the atoms are being filled with electrons. That means that it is the 4s electrons which can be thought of as being on the outside of the atom, and so determine its size. It also means that the 3d orbitals are slightly closer to the nucleus than the 4s - and so offer some screening. You will find this commented on in the page about electronic structures of ions.
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atomic and ionic radius
IONIC RADIUS Ions aren't the same size as the atoms they come from. Compare the sizes of sodium and chloride ions with the sizes of sodium and chlorine atoms.
Positive ions Positive ions are smaller than the atoms they come from. Sodium is 2,8,1; Na+ is 2,8. You've lost a whole layer of electrons, and the remaining 10 electrons are being pulled in by the full force of 11 protons. Negative ions Negative ions are bigger than the atoms they come from. Chlorine is 2,8,7; Cl- is 2,8,8. Although the electrons are still all in the 3-level, the extra repulsion produced by the incoming electron causes the atom to expand. There are still only 17 protons, but they are now having to hold 18 electrons.
Where would you like to go now? To the atomic properties menu . . . To the atomic structure and bonding menu . . . To Main Menu . . .
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atomic and ionic radius
© Jim Clark 2000 (modified 2004)
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Bonding menu
Understanding Chemistry
BONDING MENU
Ionic bonding . . . Includes a simple view of ionic bonding and the way you need to modify this for A'level purposes. Covalent bonding . . . Includes a simple view of covalent bonding (single and double) and the modifications needed for A'level purposes. Co-ordinate (dative covalent) bonding . . . Explains what co-ordinate (dative covalent) bonding is, and looks at a wide range of examples. Electronegativity . . . Explains what electronegativity is and how it varies around the Periodic Table. Describes and explains how electronegativity differences determine the type of bond formed. Looks at polar bonds and molecules. Shapes of simple molecules and ions . . . Explains how to work out the shapes of a wide range of simple molecules and ions. Metallic bonding . . . A simple explanation of the forces holding metals together. http://www.chemguide.co.uk/atoms/bondingmenu.html (1 of 2)30/12/2004 11:00:38
Bonding menu
van der Waals forces . . . A description of van der Waals forces (temporary fluctuating dipole and dipole-dipole interactions) causing attractions between individual molecules. Hydrogen bonding . . . An explanation of how hydrogen bonding arises and its effect on boiling points.
Bonding in organic compounds . . . This leads you to the bonding menu in the organic section of this site in case you are only interested in bonding in organic compounds.
Go to atomic structure and bonding menu . . . Go to Main Menu . . .
© Jim Clark 2000
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ionic (electrovalent) bonding
IONIC (ELECTROVALENT) BONDING
This page explains what ionic (electrovalent) bonding is. It starts with a simple picture of the formation of ions, and then modifies it slightly for A'level purposes.
A simple view of ionic bonding The importance of noble gas structures At a simple level (like GCSE) a lot of importance is attached to the electronic structures of noble gases like neon or argon which have eight electrons in their outer energy levels (or two in the case of helium). These noble gas structures are thought of as being in some way a "desirable" thing for an atom to have. You may well have been left with the strong impression that when other atoms react, they try to organise things such that their outer levels are either completely full or completely empty.
Note: The central role given to noble gas structures is very much an over-simplification. We shall have to spend some time later on demolishing the concept!
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ionic (electrovalent) bonding
Ionic bonding in sodium chloride Sodium (2,8,1) has 1 electron more than a stable noble gas structure (2,8). If it gave away that electron it would become more stable. Chlorine (2,8,7) has 1 electron short of a stable noble gas structure (2,8,8). If it could gain an electron from somewhere it too would become more stable. The answer is obvious. If a sodium atom gives an electron to a chlorine atom, both become more stable.
The sodium has lost an electron, so it no longer has equal numbers of electrons and protons. Because it has one more proton than electron, it has a charge of 1+. If electrons are lost from an atom, positive ions are formed. Positive ions are sometimes called cations. The chlorine has gained an electron, so it now has one more electron than proton. It therefore has a charge of 1-. If electrons are gained by an atom, negative ions are formed. A negative ion is sometimes called an anion. The nature of the bond The sodium ions and chloride ions are held together by the strong electrostatic attractions between the positive and negative charges. The formula of sodium chloride You need one sodium atom to provide the extra electron for one chlorine atom, so they combine together 1:1. The formula is therefore NaCl.
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ionic (electrovalent) bonding
Some other examples of ionic bonding magnesium oxide
Again, noble gas structures are formed, and the magnesium oxide is held together by very strong attractions between the ions. The ionic bonding is stronger than in sodium chloride because this time you have 2+ ions attracting 2- ions. The greater the charge, the greater the attraction. The formula of magnesium oxide is MgO. calcium chloride
This time you need two chlorines to use up the two outer electrons in the calcium. The formula of calcium chloride is therefore CaCl2. potassium oxide
Again, noble gas structures are formed. It takes two potassiums to supply the electrons the oxygen needs. The formula of potassium oxide is K2O. http://www.chemguide.co.uk/atoms/bonding/ionic.html (3 of 7)30/12/2004 11:00:42
ionic (electrovalent) bonding
THE A'LEVEL VIEW OF IONIC BONDING ●
●
Electrons are transferred from one atom to another resulting in the formation of positive and negative ions. The electrostatic attractions between the positive and negative ions hold the compound together.
So what's new? At heart - nothing. What needs modifying is the view that there is something magic about noble gas structures. There are far more ions which don't have noble gas structures than there are which do. Some common ions which don't have noble gas structures You may have come across some of the following ions in a basic course like GCSE. They are all perfectly stable , but not one of them has a noble gas structure. Fe3+
[Ar]3d5
Cu2+
[Ar]3d9
Zn2+
[Ar]3d10
Ag+
[Kr]4d10
Pb2+
[Xe]4f145d106s2
Noble gases (apart from helium) have an outer electronic structure ns2np6.
Note: If you aren't happy about writing electronic structures using of s, p and d notation, follow this link before you go on. Return to this page via the menus or by using the BACK button on your browser.
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ionic (electrovalent) bonding
Apart from some elements at the beginning of a transition series (scandium forming Sc3+ with an argon structure, for example), all transition elements and any metals following a transition series (like tin and lead in Group 4, for example) will have structures like those above. That means that the only elements to form positive ions with noble gas structures (apart from odd ones like scandium) are those in groups 1 and 2 of the Periodic Table and aluminium in group 3 (boron in group 3 doesn't form ions). Negative ions are tidier! Those elements in Groups 5, 6 and 7 which form simple negative ions all have noble gas structures. If elements aren't aiming for noble gas structures when they form ions, what decides how many electrons are transferred? The answer lies in the energetics of the process by which the compound is made.
Warning! From here to the bottom of this page goes beyond anything you are likely to need for A'level purposes. It is included for interest only.
What determines what the charge is on an ion? Elements combine to make the compound which is as stable as possible - the one in which the greatest amount of energy is evolved in its making. The more charges a positive ion has, the greater the attraction towards its accompanying negative ion. The greater the attraction, the more energy is released when the ions come together. That means that elements forming positive ions will tend to give away as many electrons as possible. But there's a down-side to this. Energy is needed to remove electrons from atoms. This is called ionisation energy. The more electrons you remove, the greater the total ionisation energy becomes. Eventually the total ionisation energy needed becomes so great that the energy released when the attractions are set up between positive and negative ions isn't large enough to cover it.
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The element forms the ion which makes the compound most stable - the one in which most energy is released over-all. For example, why is calcium chloride CaCl2 rather than CaCl or CaCl3? If one mole of CaCl (containing Ca+ ions) is made from its elements, it is possible to estimate that about 171 kJ of heat is evolved. However, making CaCl2 (containing Ca2+ ions) releases more heat. You get 795 kJ. That extra amount of heat evolved makes the compound more stable, which is why you get CaCl2 rather than CaCl. What about CaCl3 (containing Ca3+ ions)? To make one mole of this, you can estimate that you would have to put in 1341 kJ. This makes this compound completely non-viable. Why is so much heat needed to make CaCl3? It is because the third ionisation energy (the energy needed to remove the third electron) is extremely high (4940 kJ mol-1) because the electron is being removed from the 3-level rather than the 4-level. Because it is much closer to the nucleus than the first two electrons removed, it is going to be held much more strongly.
Note: It would pay you to read about ionisation energies if you really want to understand this. You could also go to a standard text book and investigate Born-Haber Cycles.
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A similar sort of argument applies to the negative ion. For example, oxygen forms an O2- ion rather than an O- ion or an O3- ion, because compounds containing the O2- ion turn out to be the most energetically stable.
Where would you like to go now? To explore the physical properties of ionic compounds . . . To the bonding menu . . . To the atomic structure and bonding menu . . . To Main Menu . . .
© Jim Clark 2000
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ionic structures
IONIC STRUCTURES
This page explains the relationship between the arrangement of the ions in a typical ionic solid like sodium chloride and its physical properties melting point, boiling point, brittleness, solubility and electrical behaviour.
Note: If you need to revise how ionic bonding arises, then you might like to follow this link. It isn't important for understanding this page, however.
The structure of a typical ionic solid - sodium chloride How the ions are arranged in sodium chloride Sodium chloride is taken as a typical ionic compound. Compounds like this consist of a giant (endlessly repeating) lattice of ions. So sodium chloride (and any other ionic compound) is described as having a giant ionic structure. You should be clear that giant in this context doesn't just mean very large. It means that you can't state exactly how many ions there are. There could be billions of sodium ions and chloride ions packed together, or trillions, or whatever - it simply depends how big the crystal is. That is different from, say, a water molecule which always contains exactly 2 hydrogen atoms and one oxygen atom - never more and never less. A small representative bit of a sodium chloride lattice looks like this:
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ionic structures
If you look at the diagram carefully, you will see that the sodium ions and chloride ions alternate with each other in each of the three dimensions. This diagram is easy enough to draw with a computer, but extremely difficult to draw convincingly by hand. We normally draw an "exploded" version which looks like this:
Only those ions joined by lines are actually touching each other. The sodium ion in the centre is being touched by 6 chloride ions. By chance we might just as well have centred the diagram around a chloride ion that, of course, would be touched by 6 sodium ions. Sodium chloride is described as being 6:6-co-ordinated. You must remember that this diagram represents only a tiny part of the whole sodium chloride crystal. The pattern repeats in this way over countless ions.
How to draw this structure Draw a perfect square:
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ionic structures
Now draw an identical square behind this one and offset a bit. You might have to practice a bit to get the placement of the two squares right. If you get it wrong, the ions get all tangled up with each other in your final diagram.
Turn this into a perfect cube by joining the squares together:
Now the tricky bit! Subdivide this big cube into 8 small cubes by joining the mid point of each edge to the mid point of the edge opposite it. To complete the process you will also have to join the mid point of each face (easily found once you've joined the edges) to the mid point of the opposite face.
Now all you have to do is put the ions in. Use different colours or different sizes for the two different ions, and don't forget a key. It doesn't matter whether you end up with a sodium ion or a chloride ion in the centre of http://www.chemguide.co.uk/atoms/structures/ionicstruct.html (3 of 7)30/12/2004 11:00:48
ionic structures
the cube - all that matters is that they alternate in all three dimensions.
You should be able to draw a perfectly adequate free-hand sketch of this in under two minutes - less than one minute if you're not too fussy!
Why is sodium chloride 6:6-co-ordinated? The more attraction there is between the positive and negative ions, the more energy is released. The more energy that is released, the more energetically stable the structure becomes. That means that to gain maximum stability, you need the maximum number of attractions. So why does each ion surround itself with 6 ions of the opposite charge? That represents the maximum number of chloride ions that you can fit around a central sodium ion before the chloride ions start touching each other. If they start touching, you introduce repulsions into the crystal which makes it less stable.
Note: If the positive ion is big enough, you can fit 8 chloride ions around it. For example, caesium ions are significantly bigger than sodium ions, and so caesium chloride is 8:8-coordinated. It can gain stability from the extra attractions without any problems because of repulsion due to ions with the same charge touching each other. The structure of caesium chloride isn't on any current A'level syllabuses.
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The physical properties of sodium chloride Sodium chloride has a high melting and boiling point There are strong electrostatic attractions between the positive and negative ions, and it takes a lot of heat energy to overcome them. Ionic substances all have high melting and boiling points. Differences between ionic substances will depend on things like: ●
The number of charges on the ions Magnesium oxide has exactly the same structure as sodium chloride, but a much higher melting and boiling point. The 2+ and 2- ions attract each other more strongly than 1+ attracts 1-.
●
The sizes of the ions If the ions are smaller they get closer together and so the electrostatic attractions are greater. Rubidium iodide, for example, melts and boils at slightly lower temperatures than sodium chloride, because both rubidium and iodide ions are bigger than sodium and chloride ions. The attractions are less between the bigger ions and so less heat energy is needed to separate them.
Sodium chloride crystals are brittle Brittleness is again typical of ionic substances. Imagine what happens to the crystal if a stress is applied which shifts the ion layers slightly.
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repels itself to pieces!
Sodium chloride is soluble in water Many ionic solids are soluble in water - although not all. It depends on whether there are big enough attractions between the water molecules and the ions to overcome the attractions between the ions themselves. Positive ions are attracted to the lone pairs on water molecules and coordinate (dative covalent) bonds may form. Water molecules form hydrogen bonds with negative ions.
Note: The bonding in hydrated metal ions is covered in the page on co-ordinate bonding. The bonding between negative ions like chloride ions and water molecules is covered in the page on hydrogen bonding.
Sodium chloride is insoluble in organic solvents This is also typical of ionic solids. The attractions between the solvent molecules and the ions aren't big enough to overcome the attractions holding the crystal together.
The electrical behaviour of sodium chloride Solid sodium chloride doesn't conduct electricity, because there are no electrons which are free to move. Molten sodium chloride undergoes electrolysis, which involves conduction of electricity because of the movement of the ions. In the process, sodium and chlorine are produced. This is a chemical change rather than a physical process.
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Where would you like to go now? To the structures menu . . . To the atomic structure and bonding menu . . . To Main Menu . . .
© Jim Clark 2000
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co-ordinate (dative covalent) bonding
CO-ORDINATE (DATIVE COVALENT) BONDING
This page explains what co-ordinate (also called dative covalent) bonding is. You need to have a reasonable understanding of simple covalent bonding before you start.
Important! If you are uncertain about covalent bonding follow this link before you go on with this page.
Co-ordinate (dative covalent) bonding A covalent bond is formed by two atoms sharing a pair of electrons. The atoms are held together because the electron pair is attracted by both of the nuclei. In the formation of a simple covalent bond, each atom supplies one electron to the bond - but that doesn't have to be the case. A co-ordinate bond (also called a dative covalent bond) is a covalent bond (a shared pair of electrons) in which both electrons come from the same atom. For the rest of this page, we shall use the term co-ordinate bond - but if you prefer to call it a dative covalent bond, that's not a problem!
The reaction between ammonia and hydrogen chloride If these colourless gases are allowed to mix, a thick white smoke of solid ammonium chloride is formed.
Ammonium ions, NH4+, are formed by the transfer of a hydrogen ion from the hydrogen chloride to the lone pair of electrons on the ammonia molecule. http://www.chemguide.co.uk/atoms/bonding/dative.html (1 of 9)30/12/2004 11:00:59
co-ordinate (dative covalent) bonding
When the ammonium ion, NH4+, is formed, the fourth hydrogen is attached by a dative covalent bond, because only the hydrogen's nucleus is transferred from the chlorine to the nitrogen. The hydrogen's electron is left behind on the chlorine to form a negative chloride ion. Once the ammonium ion has been formed it is impossible to tell any difference between the dative covalent and the ordinary covalent bonds. Although the electrons are shown differently in the diagram, there is no difference between them in reality. Representing co-ordinate bonds In simple diagrams, a co-ordinate bond is shown by an arrow. The arrow points from the atom donating the lone pair to the atom accepting it.
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co-ordinate (dative covalent) bonding
The H3O+ ion is variously called the hydroxonium ion, the hydronium ion or the oxonium ion. In an introductory chemistry course (such as GCSE), whenever you have talked about hydrogen ions (for example in acids), you have actually been talking about the hydroxonium ion. A raw hydrogen ion is simply a proton, and is far too reactive to exist on its own in a test tube. If you write the hydrogen ion as H+(aq), the "(aq)" represents the water molecule that the hydrogen ion is attached to. When it reacts with something (an alkali, for example), the hydrogen ion simply becomes detached from the water molecule again. Note that once the co-ordinate bond has been set up, all the hydrogens attached to the oxygen are exactly equivalent. When a hydrogen ion breaks away again, it could be any of the three.
The reaction between ammonia and boron trifluoride, BF3 If you have recently read the page on covalent bonding, you may remember boron trifluoride as a compound which doesn't have a noble gas structure around the boron atom. The boron only has 3 pairs of electrons in its bonding level, whereas there would be room for 4 pairs. BF3 is described as being electron deficient. The lone pair on the nitrogen of an ammonia molecule can be used to overcome that deficiency, and a compound is formed involving a coordinate bond. http://www.chemguide.co.uk/atoms/bonding/dative.html (3 of 9)30/12/2004 11:00:59
co-ordinate (dative covalent) bonding
Using lines to represent the bonds, this could be drawn more simply as:
The second diagram shows another way that you might find co-ordinate bonds drawn. The nitrogen end of the bond has become positive because the electron pair has moved away from the nitrogen towards the boron - which has therefore become negative. We shan't use this method again - it's more confusing than just using an arrow.
The structure of aluminium chloride Aluminium chloride sublimes (turns straight from a solid to a gas) at 178°C. If it contained ions it would have a very high melting and boiling point because of the strong attractions between the positive and negative ions. The implication is that it must be covalent. The dots-and-crosses diagram shows only the outer electrons. AlCl3, like BF3, is electron deficient. There is likely to be a similarity, because aluminium and boron are in the same group of the Periodic Table, as are fluorine and chlorine.
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co-ordinate (dative covalent) bonding
Measurements of the relative formula mass of aluminium chloride show that its formula in the solid is not AlCl3, but Al2Cl6. It exists as a dimer (two molecules joined together). The bonding between the two molecules is co-ordinate, using lone pairs on the chlorine atoms. Each chlorine atom has 3 lone pairs, but only the two important ones are shown in the line diagram.
Note: The uninteresting electrons on the chlorines have been faded in colour to make the co-ordinate bonds show up better. There's nothing special about those two particular lone pairs - they just happen to be the ones pointing in the right direction.
Energy is released when the two co-ordinate bonds are formed, and so the dimer is more stable than two separate AlCl3 molecules.
The bonding in hydrated metal ions Water molecules are strongly attracted to ions in solution - the water molecules clustering around the positive or negative ions. In many cases, the attractions are so great that formal bonds are made, and this is true of almost all positive metal ions. Ions with water molecules attached are described as hydrated ions.
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co-ordinate (dative covalent) bonding
Although aluminium chloride is covalent, when it dissolves in water, ions are produced. Six water molecules bond to the aluminium to give an ion with the formula Al(H2O)63+. It's called the hexaaquaaluminium ion which translates as six ("hexa") water molecules ("aqua") wrapped around an aluminium ion. The bonding in this (and the similar ions formed by the great majority of other metals) is co-ordinate (dative covalent) using lone pairs on the water molecules.
Aluminium is 1s22s22p63s23px1. When it forms an Al3+ ion it loses the 3level electrons to leave 1s22s22p6. That means that all the 3-level orbitals are now empty. The aluminium reorganises (hybridises) six of these (the 3s, three 3p, and two 3d) to produce six new orbitals all with the same energy. These six hybrid orbitals accept lone pairs from six water molecules. You might wonder why it chooses to use six orbitals rather than four or eight or whatever. Six is the maximum number of water molecules it is possible to fit around an aluminium ion (and most other metal ions). By making the maximum number of bonds, it releases most energy and so becomes most energetically stable.
Only one lone pair is shown on each water molecule. The other lone pair is pointing away from the aluminium and so isn't involved in the bonding. The resulting ion looks like this:
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co-ordinate (dative covalent) bonding
Because of the movement of electrons towards the centre of the ion, the 3+ charge is no longer located entirely on the aluminium, but is now spread over the whole of the ion.
Note: Dotted arrows represent lone pairs coming from water molecules behind the plane of the screen or paper. Wedge shaped arrows represent bonds from water molecules in front of the plane of the screen or paper.
Two more molecules
Note: Only one current A'level syllabus wants these two. Check yours! If you haven't got a copy of your syllabus, follow this link to find out how to get one.
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co-ordinate (dative covalent) bonding
Carbon monoxide, CO Carbon monoxide can be thought of as having two ordinary covalent bonds between the carbon and the oxygen plus a co-ordinate bond using a lone pair on the oxygen atom.
Nitric acid, HNO3 In this case, one of the oxygen atoms can be thought of as attaching to the nitrogen via a co-ordinate bond using the lone pair on the nitrogen atom.
In fact this structure is misleading because it suggests that the two oxygen atoms on the right-hand side of the diagram are joined to the nitrogen in different ways. Both bonds are actually identical in length and strength, and so the arrangement of the electrons must be identical. There is no way of showing this using a dots-and-crosses picture. The bonding involves delocalisation.
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co-ordinate (dative covalent) bonding
If you are interested: The bonding is rather similar to the bonding in the ethanoate ion (although without the negative charge). You will find thisdescribed on a page about the acidity of organic acids.
Where would you like to go now? To the bonding menu . . . To the atomic structure and bonding menu . . . To Main Menu . . .
© Jim Clark 2000
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covalent bonding - single bonds
COVALENT BONDING - SINGLE BONDS
This page explains what covalent bonding is. It starts with a simple picture of the single covalent bond, and then modifies it slightly for A'level purposes. It also takes a more sophisticated view (beyond A'level) if you are interested. You will find a link to a page on double covalent bonds at the bottom of the page.
A simple view of covalent bonding The importance of noble gas structures At a simple level (like GCSE) a lot of importance is attached to the electronic structures of noble gases like neon or argon which have eight electrons in their outer energy levels (or two in the case of helium). These noble gas structures are thought of as being in some way a "desirable" thing for an atom to have. You may well have been left with the strong impression that when other atoms react, they try to achieve noble gas structures. As well as achieving noble gas structures by transferring electrons from one atom to another as in ionic bonding, it is also possible for atoms to reach these stable structures by sharing electrons to give covalent bonds. Some very simple covalent molecules Chlorine For example, two chlorine atoms could both achieve stable structures by sharing their single unpaired electron as in the diagram.
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covalent bonding - single bonds
The fact that one chlorine has been drawn with electrons marked as crosses and the other as dots is simply to show where all the electrons come from. In reality there is no difference between them. The two chlorine atoms are said to be joined by a covalent bond. The reason that the two chlorine atoms stick together is that the shared pair of electrons is attracted to the nucleus of both chlorine atoms. Hydrogen
Hydrogen atoms only need two electrons in their outer level to reach the noble gas structure of helium. Once again, the covalent bond holds the two atoms together because the pair of electrons is attracted to both nuclei. Hydrogen chloride
The hydrogen has a helium structure, and the chlorine an argon structure.
Covalent bonding at A'level Cases where there isn't any difference from the simple view If you stick closely to modern A'level syllabuses, there is little need to http://www.chemguide.co.uk/atoms/bonding/covalent.html (2 of 13)30/12/2004 11:01:09
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move far from the simple (GCSE) view. The only thing which must be changed is the over-reliance on the concept of noble gas structures. Most of the simple molecules you draw do in fact have all their atoms with noble gas structures. For example:
Even with a more complicated molecule like PCl3, there's no problem. In this case, only the outer electrons are shown for simplicity. Each atom in this structure has inner layers of electrons of 2,8. Again, everything present has a noble gas structure.
Cases where the simple view throws up problems Boron trifluoride, BF3
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A boron atom only has 3 electrons in its outer level, and there is no possibility of it reaching a noble gas structure by simple sharing of electrons. Is this a problem? No. The boron has formed the maximum number of bonds that it can in the circumstances, and this is a perfectly valid structure. Energy is released whenever a covalent bond is formed. Because energy is being lost from the system, it becomes more stable after every covalent bond is made. It follows, therefore, that an atom will tend to make as many covalent bonds as possible. In the case of boron in BF3, three bonds is the maximum possible because boron only has 3 electrons to share.
Note: You might perhaps wonder why boron doesn't form ionic bonds with fluorine instead. Boron doesn't form ions because the total energy needed to remove three electrons to form a B3+ ion is simply too great to be recoverable when attractions are set up between the boron and fluoride ions.
Phosphorus(V) chloride, PCl5 In the case of phosphorus 5 covalent bonds are possible - as in PCl5. Phosphorus forms two chlorides - PCl3 and PCl5. When phosphorus burns in chlorine both are formed - the majority product depending on how much chlorine is available. We've already looked at the structure of PCl3. The diagram of PCl5 (like the previous diagram of PCl3) shows only the outer electrons.
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Notice that the phosphorus now has 5 pairs of electrons in the outer level - certainly not a noble gas structure. You would have been content to draw PCl3 at GCSE, but PCl5 would have looked very worrying. Why does phosphorus sometimes break away from a noble gas structure and form five bonds? In order to answer that question, we need to explore territory beyond the limits of current A'level syllabuses. Don't be put off by this! It isn't particularly difficult, and is extremely useful if you are going to understand the bonding in some important organic compounds.
A more sophisticated view of covalent bonding The bonding in methane, CH4
Warning! If you aren't happy with describing electron arrangements in s and p notation, and with the shapes of s and p orbitals, you need to read about orbitals before you go on. Use the BACK button on your browser to return quickly to this point.
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What is wrong with the dots-and-crosses picture of bonding in methane? We are starting with methane because it is the simplest case which illustrates the sort of processes involved. You will remember that the dotsand-crossed picture of methane looks like this.
There is a serious mis-match between this structure and the modern electronic structure of carbon, 1s22s22px12py1. The modern structure shows that there are only 2 unpaired electrons for hydrogens to share with, instead of the 4 which the simple view requires. You can see this more readily using the electrons-in-boxes notation. Only the 2level electrons are shown. The 1s2 electrons are too deep inside the atom to be involved in bonding. The only electrons directly available for sharing are the 2p electrons. Why then isn't methane CH2? Promotion of an electron
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When bonds are formed, energy is released and the system becomes more stable. If carbon forms 4 bonds rather than 2, twice as much energy is released and so the resulting molecule becomes even more stable. There is only a small energy gap between the 2s and 2p orbitals, and so it pays the carbon to provide a small amount of energy to promote an electron from the 2s to the empty 2p to give 4 unpaired electrons. The extra energy released when the bonds form more than compensates for the initial input.
Note: People sometimes worry that the promoted electron is drawn as an up-arrow, whereas it started as a down-arrow. It simply makes the diagram look tidier - nothing very sophisticated is going on!
Now that we've got 4 unpaired electrons ready for bonding, another problem arises. In methane all the carbon-hydrogen bonds are identical, but our electrons are in two different kinds of orbitals. You aren't going to get four identical bonds unless you start from four identical orbitals. Hybridisation The electrons rearrange themselves again in a process called hybridisation. This reorganises the electrons into four identical hybrid orbitals called sp3 hybrids (because they are made from one s orbital and three p orbitals). You should read "sp3" as "s p three" - not as "s p cubed".
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sp3 hybrid orbitals look a bit like half a p orbital, and they arrange themselves in space so that they are as far apart as possible. You can picture the nucleus as being at the centre of a tetrahedron (a triangularly based pyramid) with the orbitals pointing to the corners. For clarity, the nucleus is drawn far larger than it really is.
What happens when the bonds are formed? Remember that hydrogen's electron is in a 1s orbital - a spherically symmetric region of space surrounding the nucleus where there is some fixed chance (say 95%) of finding the electron. When a covalent bond is formed, the atomic orbitals (the orbitals in the individual atoms) merge to produce a new molecular orbital which contains the electron pair which creates the bond.
Four molecular orbitals are formed, looking rather like the original sp3 hybrids, but with a hydrogen nucleus embedded in each lobe. Each orbital holds the 2 electrons that we've previously drawn as a dot and a cross. The principles involved - promotion of electrons if necessary, then hybridisation, followed by the formation of molecular orbitals - can be applied to any covalently-bound molecule.
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Note: You will find this bit on methane repeated in the organic section of this site. That article on methane goes on to look at the formation of carbon-carbon single bonds in ethane.
The bonding in the phosphorus chlorides, PCl3 and PCl5 What's wrong with the simple view of PCl3? This diagram only shows the outer (bonding) electrons.
Nothing is wrong with this! (Although it doesn't account for the shape of the molecule properly.) If you were going to take a more modern look at it, the argument would go like this: Phosphorus has the electronic structure 1s22s22p63s23px13py13pz1. If we look only at the outer electrons as "electrons-in-boxes":
There are 3 unpaired electrons that can be used to form bonds with 3 chlorine atoms. The four 3-level orbitals hybridise to produce 4 equivalent sp3 hybrids just like in carbon - except that one of these hybrid
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orbitals contains a lone pair of electrons.
Each of the 3 chlorines then forms a covalent bond by merging the atomic orbital containing its unpaired electron with one of the phosphorus unpaired electrons to make 3 molecular orbitals. You might wonder whether all this is worth the bother! Probably not! It is worth it with PCl5, though. What's wrong with the simple view of PCl5? You will remember that the dots-and-crosses picture of PCl5 looks awkward because the phosphorus doesn't end up with a noble gas structure. This diagram also shows only the outer electrons.
In this case, a more modern view makes things look better by abandoning any pretence of worrying about noble gas structures. If the phosphorus is going to form PCl5 it has first to generate 5 unpaired electrons. It does this by promoting one of the electrons in the 3s orbital to the next available higher energy orbital. Which higher energy orbital? It uses one of the 3d orbitals. You might http://www.chemguide.co.uk/atoms/bonding/covalent.html (10 of 13)30/12/2004 11:01:09
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have expected it to use the 4s orbital because this is the orbital that fills before the 3d when atoms are being built from scratch. Not so! Apart from when you are building the atoms in the first place, the 3d always counts as the lower energy orbital.
This leaves the phosphorus with this arrangement of its electrons:
The 3-level electrons now rearrange (hybridise) themselves to give 5 hybrid orbitals, all of equal energy. They would be called sp3d hybrids because that's what they are made from.
The electrons in each of these orbitals would then share space with electrons from five chlorines to make five new molecular orbitals - and hence five covalent bonds. Why does phosphorus form these extra two bonds? It puts in an amount of energy to promote an electron, which is more than paid back when the new bonds form. Put simply, it is energetically profitable for the phosphorus to form the extra bonds. The advantage of thinking of it in this way is that it completely ignores the
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question of whether you've got a noble gas structure, and so you don't worry about it.
A non-existent compound - NCl5 Nitrogen is in the same Group of the Periodic Table as phosphorus, and you might expect it to form a similar range of compounds. In fact, it doesn't. For example, the compound NCl3 exists, but there is no such thing as NCl5. Nitrogen is 1s22s22px12py12pz1. The reason that NCl5 doesn't exist is that in order to form five bonds, the nitrogen would have to promote one of its 2s electrons. The problem is that there aren't any 2d orbitals to promote an electron into - and the energy gap to the next level (the 3s) is far too great. In this case, then, the energy released when the extra bonds are made isn't enough to compensate for the energy needed to promote an electron - and so that promotion doesn't happen. Atoms will form as many bonds as possible provided it is energetically profitable.
Where would you like to go now? To explore double covalent bonding . . . To the bonding menu . . . To the atomic structure and bonding menu . . . To Main Menu . . .
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© Jim Clark 2000
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bonding in methane - sp3 hybridisation
BONDING IN METHANE AND ETHANE
Warning! If you aren't happy with describing electron arrangements in s and p notation, and with the shapes of s and p orbitals, you really should read about orbitals. Use the BACK button on your browser to return quickly to this point.
Methane, CH4 The simple view of the bonding in methane You will be familiar with drawing methane using dots and crosses diagrams, but it is worth looking at its structure a bit more closely. There is a serious mis-match between this structure and the modern electronic structure of carbon, 1s22s22px12py1. The modern structure shows that there are only 2 unpaired electrons for hydrogens to share with, instead of the 4 which the simple view requires. You can see this more readily using the electrons-in-boxes notation. Only the 2level electrons are shown. The 1s2 electrons are too deep inside the atom to be involved in bonding. The only electrons directly available for sharing are the 2p electrons. Why then isn't methane CH2? Promotion of an electron
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bonding in methane - sp3 hybridisation
When bonds are formed, energy is released and the system becomes more stable. If carbon forms 4 bonds rather than 2, twice as much energy is released and so the resulting molecule becomes even more stable. There is only a small energy gap between the 2s and 2p orbitals, and so it pays the carbon to provide a small amount of energy to promote an electron from the 2s to the empty 2p to give 4 unpaired electrons. The extra energy released when the bonds form more than compensates for the initial input.
Note: People sometimes worry that the promoted electron is drawn as an up-arrow, whereas it started as a down-arrow. It simply makes the diagram look tidier - nothing very sophisticated is going on!
Now that we've got 4 unpaired electrons ready for bonding, another problem arises. In methane all the carbon-hydrogen bonds are identical, but our electrons are in two different kinds of orbitals. You aren't going to get four identical bonds unless you start from four identical orbitals. Hybridisation The electrons rearrange themselves again in a process called hybridisation. This reorganises the electrons into four identical hybrid orbitals called sp3 hybrids (because they are made from one s orbital and three p orbitals). You should read "sp3" as "s p three" - not as "s p cubed".
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bonding in methane - sp3 hybridisation
sp3 hybrid orbitals look a bit like half a p orbital, and they arrange themselves in space so that they are as far apart as possible. You can picture the nucleus as being at the centre of a tetrahedron (a triangularly based pyramid) with the orbitals pointing to the corners. For clarity, the nucleus is drawn far larger than it really is.
What happens when the bonds are formed? Remember that hydrogen's electron is in a 1s orbital - a spherically symmetric region of space surrounding the nucleus where there is some fixed chance (say 95%) of finding the electron. When a covalent bond is formed, the atomic orbitals (the orbitals in the individual atoms) merge to produce a new molecular orbital which contains the electron pair which creates the bond.
Four molecular orbitals are formed, looking rather like the original sp3 hybrids, but with a hydrogen nucleus embedded in each lobe. Each orbital holds the 2 electrons that we've previously drawn as a dot and a cross. The principles involved - promotion of electrons if necessary, then hybridisation, followed by the formation of molecular orbitals - can be applied to any covalently-bound molecule.
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bonding in methane - sp3 hybridisation
Ethane, C2H6 The formation of molecular orbitals in ethane Ethane isn't particularly important in its own right, but is included because it is a simple example of how a carbon-carbon single bond is formed. Each carbon atom in the ethane promotes an electron and then forms sp3 hybrids exactly as we've described in methane. So just before bonding, the atoms look like this:
The hydrogens bond with the two carbons to produce molecular orbitals just as they did with methane. The two carbon atoms bond by merging their remaining sp3 hybrid orbitals end-to-end to make a new molecular orbital. The bond formed by this end-to-end overlap is called a sigma bond. The bonds between the carbons and hydrogens are also sigma bonds.
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In any sigma bond, the most likely place to find the pair of electrons is on a line between the two nuclei. Free rotation about the carbon-carbon single bond The two ends of this molecule can spin quite freely about the sigma bond so that there are, in a sense, an infinite number of possibilities for the shape of an ethane molecule. Some possible shapes are:
In each case, the left hand CH3 group has been kept in a constant position so that you can see the effect of spinning the right hand one. Other alkanes All other alkanes will be bonded in the same way: ●
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The carbon atoms will each promote an electron and then hybridise to give sp3 hybrid orbitals. The carbon atoms will join to each other by forming sigma bonds by the end-to-end overlap of their sp3 hybrid orbitals. Hydrogen atoms will join on wherever they are needed by overlapping their 1s1 orbitals with sp3 hybrid orbitals on the carbon atoms.
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Where would you like to go now? To the organic bonding menu. . . To menu of basic organic chemistry. . . To Main Menu . . .
© Jim Clark 2000
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electronic structure and atomic orbitals
ELECTRONIC STRUCTURE AND ATOMIC ORBITALS
A simple view In any introductory chemistry course you will have come across the electronic structures of hydrogen and carbon drawn as:
Note: There are many places where you could still make use of this model of the atom at A' level. It is, however, a simplification and can be misleading. It gives the impression that the electrons are circling the nucleus in orbits like planets around the sun. As you will see in a moment, it is impossible to know exactly how they are actually moving.
The circles show energy levels - representing increasing distances from the nucleus. You could straighten the circles out and draw the electronic structure as a simple energy diagram.
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electronic structure and atomic orbitals
Atomic orbitals Orbits and orbitals sound similar, but they have quite different meanings. It is essential that you understand the difference between them. The impossibility of drawing orbits for electrons To plot a path for something you need to know exactly where the object is and be able to work out exactly where it's going to be an instant later. You can't do this for electrons.
Note: In order to plot a plane's course, it is no use knowing its exact location in mid-Atlantic if you don't know its direction or speed. Equally it's no use knowing that it is travelling at 500 mph due west if you have no idea whether it is near Iceland or the Azores at that particular moment.
The Heisenberg Uncertainty Principle (not required at A'level) says loosely - that you can't know with certainty both where an electron is and where it's going next. That makes it impossible to plot an orbit for an electron around a nucleus. Is this a big problem? No. If something is impossible, you have to accept it and find a way around it. Hydrogen's electron - the 1s orbital
Note: In this diagram (and the orbital diagrams that follow), the nucleus is shown very much larger than it really is. This is just for clarity.
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electronic structure and atomic orbitals
Suppose you had a single hydrogen atom and at a particular instant plotted the position of the one electron. Soon afterwards, you do the same thing, and find that it is in a new position. You have no idea how it got from the first place to the second. You keep on doing this over and over again, and gradually build up a sort of 3D map of the places that the electron is likely to be found. In the hydrogen case, the electron can be found anywhere within a spherical space surrounding the nucleus. The diagram shows a crosssection through this spherical space. 95% of the time (or any other percentage you choose), the electron will be found within a fairly easily defined region of space quite close to the nucleus. Such a region of space is called an orbital. You can think of an orbital as being the region of space in which the electron lives.
Note: If you wanted to be absolutely 100% sure of where the electron is, you would have to draw an orbital the size of the Universe!
What is the electron doing in the orbital? We don't know, we can't know, and so we just ignore the problem! All you can say is that if an electron is in a particular orbital it will have a particular definable energy. Each orbital has a name. The orbital occupied by the hydrogen electron is called a 1s orbital. The "1" represents the fact that the orbital is in the energy level closest to the nucleus. The "s" tells you about the shape of the orbital. s orbitals are spherically symmetric around the nucleus - in each case, like a hollow ball made of rather chunky material with the nucleus at its centre.
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electronic structure and atomic orbitals
The orbital on the left is a 2s orbital. This is similar to a 1s orbital except that the region where there is the greatest chance of finding the electron is further from the nucleus - this is an orbital at the second energy level. If you look carefully, you will notice that there is another region of slightly higher electron density (where the dots are thicker) nearer the nucleus. ("Electron density" is another way of talking about how likely you are to find an electron at a particular place.) 2s (and 3s, 4s, etc) electrons spend some of their time closer to the nucleus than you might expect. The effect of this is to slightly reduce the energy of electrons in s orbitals. The nearer the nucleus the electrons get, the lower their energy. 3s, 4s (etc) orbitals get progressively further from the nucleus. p orbitals Not all electrons inhabit s orbitals (in fact, very few electrons live in s orbitals). At the first energy level, the only orbital available to electrons is the 1s orbital, but at the second level, as well as a 2s orbital, there are also orbitals called 2p orbitals. A p orbital is rather like 2 identical balloons tied together at the nucleus. The diagram on the right is a cross-section through that 3-dimensional region of space. Once again, the orbital shows where there is a 95% chance of finding a particular electron.
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electronic structure and atomic orbitals
Beyond A'level: If you imagine a horizontal plane through the nucleus, with one lobe of the orbital above the plane and the other beneath it, there is a zero probability of finding the electron on that plane. So how does the electron get from one lobe to the other if it can never pass through the plane of the nucleus? For A'level chemistry you just have to accept that it does! If you want to find out more, read about the wave nature of electrons.
Unlike an s orbital, a p orbital points in a particular direction - the one drawn points up and down the page. At any one energy level it is possible to have three absolutely equivalent p orbitals pointing mutually at right angles to each other. These are arbitrarily given the symbols px, py and pz. This is simply for convenience - what you might think of as the x, y or z direction changes constantly as the atom tumbles in space. The p orbitals at the second energy level are called 2px, 2py and 2pz. There are similar orbitals at subsequent levels - 3px, 3py, 3pz, 4px, 4py, 4pz and so on. All levels except for the first level have p orbitals. At the higher levels the lobes get more elongated, with the most likely place to find the electron more distant from the nucleus.
Fitting electrons into orbitals Because for the moment we are only interested in the electronic structures of hydrogen and carbon, we don't need to concern ourselves with what happens beyond the second energy level. Remember:
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electronic structure and atomic orbitals
At the first level there is only one orbital - the 1s orbital. At the second level there are four orbitals - the 2s, 2px, 2py and 2pz orbitals. Each orbital can hold either 1 or 2 electrons, but no more. "Electrons-in-boxes" Orbitals can be represented as boxes with the electrons in them shown as arrows. Often an up-arrow and a down-arrow are used to show that the electrons are in some way different.
Beyond A'level: The need to have all electrons in an atom different comes out of quantum theory. If they live in different orbitals, that's fine - but if they are both in the same orbital there has to be some subtle distinction between them. Quantum theory allocates them a property known as "spin" which is what the arrows are intended to suggest.
A 1s orbital holding 2 electrons would be drawn as shown on the right, but it can be written even more quickly as 1s2. This is read as "one s two" - not as "one s squared". You mustn't confuse the two numbers in this notation:
The order of filling orbitals Electrons fill low energy orbitals (closer to the nucleus) before they fill higher energy ones. Where there is a choice between orbitals of equal energy, they fill the orbitals singly as far as possible. The diagram (not to scale) summarises the energies of the various orbitals in the first and second levels. http://www.chemguide.co.uk/basicorg/bonding/orbitals.html (6 of 8)30/12/2004 11:01:21
electronic structure and atomic orbitals
Notice that the 2s orbital has a slightly lower energy than the 2p orbitals. That means that the 2s orbital will fill with electrons before the 2p orbitals. All the 2p orbitals have exactly the same energy.
The electronic structure of hydrogen Hydrogen only has one electron and that will go into the orbital with the lowest energy - the 1s orbital. Hydrogen has an electronic structure of 1s1. We have already described this orbital earlier.
The electronic structure of carbon Carbon has six electrons. Two of them will be found in the 1s orbital close to the nucleus. The next two will go into the 2s orbital. The remaining ones will be in two separate 2p orbitals. This is because the p orbitals all have the same energy and the electrons prefer to be on their own if that's the case.
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electronic structure and atomic orbitals
Note: People sometimes wonder why the electrons choose to go into the 2px and 2py orbitals rather than the 2pz. They don't! All of the 2p orbitals are exactly equivalent, and the names we give them are entirely arbitrary. It just looks tidier if we call the orbitals the electrons occupy the 2px and 2py.
The electronic structure of carbon is normally written 1s22s22px12py1.
Where would you like to go now? To the organic bonding menu. . . To menu of basic organic chemistry. . . To Main Menu . . .
© Jim Clark 2000
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Organic Bonding Menu
Understanding Chemistry
ORGANIC BONDING MENU
Electronic structure and orbitals . . . An introduction to the arrangement of electrons in atoms - leading to the modern electronic structures of carbon and hydrogen. Bonding in methane . . . Covers bonding in methane and ethane, including a simple look at hybridisation. Bonding in ethene . . . Covers bonding in ethene, including a simple look at hybridisation. Bonding in benzene - the Kekulé structure . . . A description of the Kekulé structure for benzene and the reasons (including hydrogenation energies) why it isn't satisfactory. Bonding in benzene - a modern orbital view . . . Covers a modern view of the bonding in benzene, including a simple look at hybridisation. Bonding in carbonyl compounds . . . Describes the carbon-oxygen double bond in methanal (including a simple look at hybridisation), but applies equally to other aldehydes and ketones like ethanal and propanone.
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Organic Bonding Menu
Electronegativity . . . An introduction to electronegativity as it applies to organic chemistry, including its causes. Bond polarity.
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© Jim Clark 2000
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Understanding Chemistry - Basic Organic Chemistry Menu
Understanding Chemistry
BASIC ORGANIC CHEMISTRY MENU
Bonding in organic compounds . . . Includes basic electronic structure, bonding in methane, ethene, benzene and carbonyl compounds, and ideas about electronegativity and bond polarity. Organic chemistry conventions . . . Includes how to name and draw organic compounds, and the use of curly arrows in reaction mechanisms. Isomerism in organic compounds . . . Includes structural isomerism and stereoisomerism (both geometric and optical). Organic acids and bases . . . Includes the acid strengths of carboxylic acids, phenols and alcohols, and the base strengths of primary amines.
Go to Main Menu . . .
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Understanding Chemistry - Basic Organic Chemistry Menu
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Organic Conventions Menu
Understanding Chemistry
ORGANIC CONVENTIONS MENU
How to draw organic molecules . . . Explains the various conventions used in drawing organic molecules. An introduction to naming organic molecules . . . A guide to understanding the names of organic compounds, including alkanes, cycloalkanes,alkenes, simple halogen compounds, alcohols, aldehydes and ketones. More organic names . . . Explains the naming of carboxylic acids and their salts, esters, acyl chlorides, acid anhydrides, amides, nitriles, amines and amino acids. It assumes that you have already read the introductory page. Naming aromatic compounds . . . Looks at the special problems involved in naming compounds containing benzene rings. It assumes that you are familiar with the naming of simple chain compounds. The use of curly arrows . . . How to use curly arrows to show the movement of electron pairs or single electrons in reaction mechanisms.
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Organic Conventions Menu
If you are interested in testing your ability to write names for organic compounds, you might like to explore these links to Dr Phil Brown's website: Multiple choice tests on organic names . . . "Type in the name" tests on organic names . . .
Go to menu of basic organic chemistry. . . Go to Main Menu . . .
© Jim Clark 2000
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How to draw organic molecules
DRAWING ORGANIC MOLECULES
This page explains the various ways that organic molecules can be represented on paper or on screen - including molecular formulae, and various forms of structural formulae.
Molecular formulae A molecular formula simply counts the numbers of each sort of atom present in the molecule, but tells you nothing about the way they are joined together. For example, the molecular formula of butane is C4H10, and the molecular formula of ethanol is C2H6O. Molecular formulae are very rarely used in organic chemistry, because they don't give any useful information about the bonding in the molecule. About the only place where you might come across them is in equations for the combustion of simple hydrocarbons, for example:
In cases like this, the bonding in the organic molecule isn't important.
Structural formulae A structural formula shows how the various atoms are bonded. There are various ways of drawing this and you will need to be familiar with all of them. Displayed formulae A displayed formula shows all the bonds in the molecule as individual
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How to draw organic molecules
lines. You need to remember that each line represents a pair of shared electrons. For example, this is a model of methane together with its displayed formula:
Notice that the way the methane is drawn bears no resemblance to the actual shape of the molecule. Methane isn't flat with 90° bond angles. This mismatch between what you draw and what the molecule actually looks like can lead to problems if you aren't careful. For example, consider the simple molecule with the molecular formula CH2Cl2. You might think that there were two different ways of arranging these atoms if you drew a displayed formula.
The chlorines could be opposite each other or at right angles to each other. But these two structures are actually exactly the same. Look at how they appear as models.
One structure is in reality a simple rotation of the other one.
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How to draw organic molecules
Note: This is all much easier to understand if you have actually got some models to play with. If your school or college hasn't given you the opportunity to play around with molecular models in the early stages of your organic chemistry course, you might consider getting hold of a cheap set. The models made by molymod are both cheap and easy to use. An introductory organic set is more than adequate. Find them at www.molymod.com. Alternatively , get hold of some coloured Plasticene and some used matches and make your own. It's cheaper, but distinctly messier!
Consider a slightly more complicated molecule, C2H5Cl. The displayed formula could be written as either of these:
But, again these are exactly the same. Look at the models.
The commonest way to draw structural formulae For anything other than the most simple molecules, drawing a fully displayed formula is a bit of a bother - especially all the carbon-hydrogen bonds. You can simplify the formula by writing, for example, CH3 or CH2 http://www.chemguide.co.uk/basicorg/conventions/draw.html (3 of 10)30/12/2004 11:01:38
How to draw organic molecules
instead of showing all these bonds. So for example, ethanoic acid would be shown in a fully displayed form and a simplified form as:
You could even condense it further to CH3COOH, and would probably do this if you had to write a simple chemical equation involving ethanoic acid. You do, however, lose something by condensing the acid group in this way, because you can't immediately see how the bonding works. You still have to be careful in drawing structures in this way. Remember from above that these two structures both represent the same molecule:
The next three structures all represent butane.
All of these are just versions of four carbon atoms joined up in a line. The only difference is that there has been some rotation about some of the carbon-carbon bonds. You can see this in a couple of models.
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How to draw organic molecules
Not one of the structural formulae accurately represents the shape of butane. The convention is that we draw it with all the carbon atoms in a straight line - as in the first of the structures above. This is even more important when you start to have branched chains of carbon atoms. The following structures again all represent the same molecule - 2-methylbutane.
The two structures on the left are fairly obviously the same - all we've done is flip the molecule over. The other one isn't so obvious until you look at the structure in detail. There are four carbons joined up in a row, with a CH3 group attached to the next-to-end one. That's exactly the same as the other two structures. If you had a model, the only difference between these three diagrams is that you have rotated some of the bonds and turned the model around a bit. To overcome this possible confusion, the convention is that you always look for the longest possible chain of carbon atoms, and then draw it horizontally. Anything else is simply hung off that chain. It doesn't matter in the least whether you draw any side groups pointing up or down. All of the following represent exactly the same molecule.
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How to draw organic molecules
If you made a model of one of them, you could turn it into any other one simply by rotating one or more of the carbon-carbon bonds. How to draw structural formulae in 3-dimensions There are occasions when it is important to be able to show the precise 3D arrangement in parts of some molecules. To do this, the bonds are shown using conventional symbols:
For example, you might want to show the 3-D arrangement of the groups around the carbon which has the -OH group in butan-2-ol. Butan-2-ol has the structural formula:
Using conventional bond notation, you could draw it as, for example:
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How to draw organic molecules
The only difference between these is a slight rotation of the bond between the centre two carbon atoms. This is shown in the two models below. Look carefully at them - particularly at what has happened to the lone hydrogen atom. In the left-hand model, it is tucked behind the carbon atom. In the right-hand model, it is in the same plane. The change is very slight.
It doesn't matter in the least which of the two arrangements you draw. You could easily invent other ones as well. Choose one of them and get into the habit of drawing 3-dimensional structures that way. My own habit (used elsewhere on this site) is to draw two bonds going back into the paper and one coming out - as in the left-hand diagram above. Notice that no attempt was made to show the whole molecule in 3dimensions in the structural formula diagrams. The CH2CH3 group was left in a simple form. Keep diagrams simple - trying to show too much detail makes the whole thing amazingly difficult to understand! Skeletal formulae In a skeletal formula, all the hydrogen atoms are removed from carbon chains, leaving just a carbon skeleton with functional groups attached to it. For example, we've just been talking about butan-2-ol. The normal structural formula and the skeletal formula look like this:
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How to draw organic molecules
In a skeletal diagram of this sort ●
●
there is a carbon atom at each junction between bonds in a chain and at the end of each bond (unless there is something else there already - like the -OH group in the example); there are enough hydrogen atoms attached to each carbon to make the total number of bonds on that carbon up to 4.
Beware! Diagrams of this sort take practice to interpret correctly - and may well not be acceptable to your examiners (see below). There are, however, some very common cases where they are frequently used. These cases involve rings of carbon atoms which are surprisingly awkward to draw tidily in a normal structural formula. Cyclohexane, C6H12, is a ring of carbon atoms each with two hydrogens attached. This is what it looks like in both a structural formula and a skeletal formula.
And this is cyclohexene, which is similar but contains a double bond:
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How to draw organic molecules
But the commonest of all is the benzene ring, C6H6, which has a special symbol of its own.
Note: Explaining exactly what this structure means needs more space than is available here. It is explained in full in two pages on the structure of benzene elsewhere in this site. It would probably be better not to follow this link unless you are actively interested in benzene chemistry at the moment - it will lead you off into quite deep water!
Deciding which sort of formula to use There's no easy, all-embracing answer to this problem. It depends more than anything else on experience - a feeling that a particular way of writing a formula is best for the situation you are dealing with. Don't worry about this - as you do more and more organic chemistry, you will probably find it will come naturally. You'll get so used to writing formulae in reaction mechanisms, or for the structures for isomers, or in simple chemical equations, that you won't even think about it. There are, however, a few guidelines that you should follow. What does your syllabus say? Different examiners will have different preferences. Check first with your syllabus. If you've down-loaded a copy of your syllabus from your Exam Board's web site, it is easy to check what they say they want. Use the "find" function on your Adobe Acrobat Reader to search the organic section(s) of the syllabus for the word "formula". You should also check recent exam papers and (particulary) mark schemes to find out what sort of formula the examiners really prefer in http://www.chemguide.co.uk/basicorg/conventions/draw.html (9 of 10)30/12/2004 11:01:38
How to draw organic molecules
given situations. You could also look at any support material published by your Board.
Note: If you haven't got a copy of your syllabus and recent exam papers, follow this link to find out how to get them.
What if you still aren't sure? Draw the most detailed formula that you can fit into the space available. If in doubt, draw a fully displayed formula. You would never lose marks for giving too much detail. Apart from the most trivial cases (for example, burning hydrocarbons), never use a molecular formula. Always show the detail around the important part(s) of a molecule. For example, the important part of an ethene molecule is the carbon-carbon double bond - so write (at the very least) CH2=CH2 and not C2H4. Where a particular way of drawing a structure is important, this will always be pointed out where it arises elsewhere on this site.
Where would you like to go now? To the organic conventions menu. . . To menu of basic organic chemistry. . . To Main Menu . . .
© Jim Clark 2000
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bonding in benzene - the Kekulé structure
BONDING IN BENZENE The Kekulé structure for benzene, C6H6 What is the Kekulé structure? Kekulé was the first to suggest a sensible structure for benzene. The carbons are arranged in a hexagon, and he suggested alternating double and single bonds between them. Each carbon atom has a hydrogen attached to it. This diagram is often simplified by leaving out all the carbon and hydrogen atoms!
In diagrams of this sort, there is a carbon atom at each corner. You have to count the bonds leaving each carbon to work out how many hydrogens there are attached to it. In this case, each carbon has three bonds leaving it. Because carbon atoms form four bonds, that means you are a bond missing - and that must be attached to a hydrogen atom. Problems with the Kekulé structure Although the Kekulé structure was a good attempt in its time, there are serious problems with it . . . Problems with the chemistry Because of the three double bonds, you might expect benzene to have reactions like ethene - only more so! Ethene undergoes addition reactions in which one of the two bonds joining the carbon atoms breaks, and the electrons are used to bond with additional atoms.
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bonding in benzene - the Kekulé structure
Benzene rarely does this. Instead, it usually undergoes substitution reactions in which one of the hydrogen atoms is replaced by something new.
Note: Follow these links to get details about the addition reactions of ethene, or the substitution reactions of benzene.
Problems with the shape Benzene is a planar molecule (all the atoms lie in one plane), and that would also be true of the Kekulé structure. The problem is that C-C single and double bonds are different lengths. C-C C=C
0.154 nm 0.134 nm
Note: "nm" means "nanometre", which is 10-9 metre.
That would mean that the hexagon would be irregular if it had the Kekulé structure, with alternating shorter and longer sides. In real benzene all the bonds are exactly the same - intermediate in length between C-C and C=C at 0.139 nm. Real benzene is a perfectly regular hexagon. Problems with the stability of benzene Real benzene is a lot more stable than the Kekulé structure would give it credit for. Every time you do a thermochemistry calculation based on the Kekulé structure, you get an answer which is wrong by about 150 kJ mol1. This is most easily shown using enthalpy changes of hydrogenation.
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bonding in benzene - the Kekulé structure
Help! It doesn't matter whether you've done any thermochemistry sums recently or not. This is all so simple that you could understand it even if you had never done any!
Hydrogenation is the addition of hydrogen to something. If, for example, you hydrogenate ethene you get ethane: CH2=CH2 + H2
CH3CH3
In order to do a fair comparison with benzene (a ring structure) we're going to compare it with cyclohexene. Cyclohexene, C6H10, is a ring of six carbon atoms containing just one C=C.
Note: If you are a bit shaky on names: cyclohexene: hex means six carbons, cyclo means in a ring, ene means with a C=C bond.
When hydrogen is added to this, cyclohexane, C6H12, is formed. The "CH" groups become CH2 and the double bond is replaced by a single one.
Note: cyclohexane: six carbons in a ring, but the ane ending means NO C=C bond.
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bonding in benzene - the Kekulé structure
The structures of cyclohexene and cyclohexane are usually simplified in the same way that the Kekulé structure for benzene is simplified - by leaving out all the carbons and hydrogens.
In the cyclohexane case, for example, there is a carbon atom at each corner, and enough hydrogens to make the total bonds on each carbon atom up to four. In this case, then, each corner represents CH2. The hydrogenation equation could be written:
The enthalpy change during this reaction is -120 kJ mol-1. In other words, when 1 mole of cyclohexene reacts, 120 kJ of heat energy is evolved.
Help! "Enthalpy change" can be translated as "heat evolved or absorbed". The negative sign shows that heat is evolved.
Where does this heat energy come from? When the reaction happens, bonds are broken (C=C and H-H) and this costs energy. Other bonds have to be made, and this releases energy. Because the bonds made are stronger than those broken, more energy is released than was used to break the original bonds and so there is a net evolution of heat energy. If the ring had two double bonds in it initially (cyclohexa-1,3-diene), exactly twice as many bonds would have to be broken and exactly twice as many made. In other words, you would expect the enthalpy change of hydrogenation of cyclohexa-1,3-diene to be exactly twice that of http://www.chemguide.co.uk/basicorg/bonding/benzene1.html (4 of 7)30/12/2004 11:01:44
bonding in benzene - the Kekulé structure
cyclohexene - that is, -240 kJ mol-1.
Note: The name (cyclohexa-1,3-diene) is unimportant. Don't worry about it unless you want to!
In fact, the enthalpy change is -232 kJ mol-1 - which isn't far off what we are predicting.
Note: Thermochemistry sums often throw up discrepancies of this sort of magnitude, and you couldn't be sure whether there was any significance in it.
Applying the same argument to the Kekulé structure for benzene (what might be called cyclohexa-1,3,5-triene), you would expect an enthalpy change of -360 kJ mol-1, because there are exactly three times as many bonds being broken and made as in the cyclohexene case.
In fact what you get is -208 kJ mol-1 - not even within distance of the predicted value! This is very much easier to see on an enthalpy diagram. Notice that in each case heat energy is released, and in each case the product is the same (cyclohexane). That means that all the reactions "fall down" to the same end point.
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bonding in benzene - the Kekulé structure
Heavy lines, solid arrows and bold numbers represent real changes. Predicted changes are shown by dotted lines and italics. The most important point to notice is that real benzene is much lower down the diagram than the Kekulé form predicts. The lower down a substance is, the more energetically stable it is. This means that real benzene is about 150 kJ mol-1 more stable than the Kekulé structure gives it credit for. This increase in stability of benzene is known as the delocalisation energy or resonance energy of benzene. The first term (delocalisation energy) is the more commonly used.
Note: If you look at the diagram closely, you will see that cyclohexa-1,3-diene is also a shade more stable than expected. There is a tiny amount of delocalisation energy involved here as well.
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bonding in benzene - the Kekulé structure
Why is benzene so much more stable than the Kekulé structure suggests? To explain that needs a separate article! Follow the first link below.
Where would you like to go now? To read about the modern view of the structure of benzene. . . To the organic bonding menu. . . To menu of basic organic chemistry. . . To Main Menu . . .
© Jim Clark 2000
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electrophilic addition reactions menu
Understanding Chemistry
ELECTROPHILIC ADDITION MECHANISMS MENU
Addition to symmetrical alkenes Covers addition to symmetrical alkenes like ethene and cyclohexene. A symmetrical alkene has the same groups attached to both ends of the carbon-carbon double bond. What is electrophilic addition? . . . An explanation of the terms addition and electrophile, together with a general mechanism for these reactions. The reaction with hydrogen halides . . . The mechanism for the reaction between ethene (and cyclohexene) and hydrogen halides (like hydrogen bromide). The reaction with sulphuric acid . . . The mechanism for the reaction between ethene (and cyclohexene) and sulphuric acid. The reaction with bromine . . . The mechanism for the reaction between ethene (and cyclohexene) and bromine.
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electrophilic addition reactions menu
Covers addition to unsymmetrical alkenes like propene. An unsymmetrical alkene has different groups attached to each end of the carbon-carbon double bond.
Warning! Don't even think about reading articles in this section until you are sure you understand the corresponding reaction(s) above!
Carbocations (carbonium ions) and their stability . . . Essential pre-reading before you tackle anything else in this section. Why unsymmetric alkenes are a problem . . . Explains the reasons behind Markovnikov's Rule, and gives a general mechanism for these more awkward reactions. This is also essential reading before you look at specific reactions. The reaction with hydrogen halides . . . The mechanism for the reaction between propene and hydrogen halides (like hydrogen bromide). The reaction with sulphuric acid . . . The mechanism for the reaction between propene and sulphuric acid. The reaction with bromine . . . The mechanism for the reaction between propene and bromine.
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electrophilic addition reactions menu
Go to menu of other types of mechanism. . . Go to Main Menu . . .
You might also be interested in: properties and reactions of alkenes . . . A survey of all the physical and chemical properties of alkenes required by UK A level syllabuses.
© Jim Clark 2000 (modified 2004)
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What is electrophilic addition?
ELECTROPHILIC ADDITION
Background Electrophilic addition happens in many of the reactions of compounds containing carbon-carbon double bonds - the alkenes. The structure of ethene We are going to start by looking at ethene, because it is the simplest molecule containing a carbon-carbon double bond. What is true of C=C in ethene will be equally true of C=C in more complicated alkenes. Ethene, C2H4, is often modelled as shown on the right. The double bond between the carbon atoms is, of course, two pairs of shared electrons. What the diagram doesn't show is that the two pairs aren't the same as each other. One of the pairs of electrons is held on the line between the two carbon nuclei as you would expect, but the other is held in a molecular orbital above and below the plane of the molecule. A molecular orbital is a region of space within the molecule where there is a high probability of finding a particular pair of electrons. In this diagram, the line between the two carbon atoms represents a normal bond - the pair of shared electrons lies in a molecular orbital on the line between the two nuclei where you would expect them to be. This sort of bond is called a sigma bond. The other pair of electrons is found somewhere in the shaded part above and below the plane of the molecule. This bond is called a pi bond. The electrons in the pi bond are free to move around anywhere in this shaded region and can move freely from one half to the other.
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What is electrophilic addition?
Note: This diagram shows a side view of an ethene molecule. The dotted lines to two of the hydrogens show bonds going back into the screen or paper away from you. The wedge shapes show bonds coming out towards you.
The pi electrons are not as fully under the control of the carbon nuclei as the electrons in the sigma bond and, because they lie exposed above and below the rest of the molecule, they are relatively open to attack by other things.
Note: Check your syllabus to see if you need to know how a pi bond is formed. Haven't got a syllabus? Find out how to get one by following this link. If you do need to know about the bonding in ethene in detail, follow this link as well.
Electrophiles An electrophile is something which is attracted to electron-rich regions in other molecules or ions. Because it is attracted to a negative region, an electrophile must be something which carries either a full positive charge, or has a slight positive charge on it somewhere.
Note: The ending ". . phile" means a liking for. For example, a francophile is someone who likes the French; an anglophile is someone who likes the English.
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What is electrophilic addition?
Ethene and the other alkenes are attacked by electrophiles. The electrophile is normally the slightly positive ( +) end of a molecule like hydrogen bromide, HBr.
Note: If you aren't sure about why some bonds are polar, read the page on electronegativity. Use the BACK button on your browser to return to this page.
Electrophiles are strongly attracted to the exposed electrons in the pi bond and reactions happen because of that initial attraction - as you will see shortly. You might wonder why fully positive ions like sodium, Na+, don't react with ethene. Although these ions may well be attracted to the pi bond, there is no possibility of the process going any further to form bonds between sodium and carbon, because sodium forms ionic bonds, whereas carbon normally forms covalent ones. Addition reactions In a sense, the pi bond is an unnecessary bond. The structure would hold together perfectly well with a single bond rather than a double bond. The pi bond often breaks and the electrons in it are used to join other atoms (or groups of atoms) onto the ethene molecule. In other words, ethene undergoes addition reactions. For example, using a general molecule X-Y . . .
Summary: electrophilic addition reactions
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What is electrophilic addition?
An addition reaction is a reaction in which two molecules join together to make a bigger one. Nothing is lost in the process. All the atoms in the original molecules are found in the bigger one. An electrophilic addition reaction is an addition reaction which happens because what we think of as the "important" molecule is attacked by an electrophile. The "important" molecule has a region of high electron density which is attacked by something carrying some degree of positive charge.
Note: When we talk about reactions of alkenes like ethene, we think of the ethene as being attacked by other molecules such as hydrogen bromide. Because ethene is the molecule we are focusing on, we quite arbitrarily think of it as the central molecule and hydrogen bromide as its attacker. There's no real justification for this, of course, apart from the fact that it helps to put things in some sort of logical pattern. In reality, the molecules just collide and may react if they have enough energy and if they are lined up correctly.
Understanding the electrophilic addition mechanism The mechanism for the reaction between ethene and a molecule X-Y It is very unlikely that any two different atoms joined together will have the same electronegativity. We are going to assume that Y is more electronegative than X, so that the pair of electrons is pulled slightly towards the Y end of the bond. That means that the X atom carries a slight positive charge.
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Note: Once again, if you aren't sure about electronegativity and bond polarity follow this link before you read on. Use the BACK button on your browser to return to this page.
The slightly positive X atom is an electrophile and is attracted to the exposed pi bond in the ethene. Now imagine what happens as they approach each other.
You are now much more likely to find the electrons in the half of the pi bond nearest the XY. As the process continues, the two electrons in the pi bond move even further towards the X until a covalent bond is made. The electrons in the X-Y bond are pushed entirely onto the Y to give a negative Y- ion.
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Help! Why does the carbon atom have a positive charge? The pi bond was originally made using an electron from each carbon atom, but both of these electrons have now been used to make a bond to the X atom. This leaves the righthand carbon atom an electron short - hence positively charged. Note also that we are only showing one of the pairs of electrons around the Y- ion. There will be other lone pairs as well, but we are only actually interested in the one we've drawn.
Important term An ion in which the positive charge is carried on a carbon atom is called a carbocation or a carbonium ion (an older term).
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In the final stage of the reaction the electrons in the lone pair on the Y- ion are strongly attracted towards the positive carbon atom. They move towards it and form a co-ordinate (dative covalent) bond between the Y and the carbon.
Help! A co-ordinate (dative covalent) bond is simply a covalent bond in which both shared electrons originate from the same atom. The bond formed between the X and the other carbon atom was also a co-ordinate bond. Once a coordinate bond has been formed there is no difference whatsoever between it and any other covalent bond.
How to write this mechanism in an exam The movements of the various electron pairs are shown using curly arrows.
Help! If you aren't sure about the use of curly arrows in mechanisms, you must follow this link before you go on. Use the BACK button on your browser to return to this page.
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Don't leave this page until you are sure that you understand how this relates to the electron pair movements drawn in the previous diagrams.
Where would you like to go now? To menu of electrophilic addition reactions. . . To menu of other types of mechanism. . . To Main Menu . . .
© Jim Clark 2000
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bonding in ethene - sp2 hybridisation
BONDING IN ETHENE
Important! You will find this much easier to understand if you first read the article about the bonding in methane. You may also find it useful to read the article on orbitals if you aren't sure about simple orbital theory.
Ethene, C2H4 The simple view of the bonding in ethene At a simple level, you will have drawn ethene showing two bonds between the carbon atoms. Each line in this diagram represents one pair of shared electrons. Ethene is actually much more interesting than this. An orbital view of the bonding in ethene Ethene is built from hydrogen atoms (1s1) and carbon atoms (1s22s22px12py1). The carbon atom doesn't have enough unpaired electrons to form the required number of bonds, so it needs to promote one of the 2s2 pair into the empty 2pz orbital. This is exactly the same as happens whenever carbon forms bonds - whatever else it ends up joined to.
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bonding in ethene - sp2 hybridisation
Important! If this isn't really clear to you, you must go and read the article about the bonding in methane.
Now there's a difference, because each carbon is only joining to three other atoms rather than four - as in methane or ethane. When the carbon atoms hybridise their outer orbitals before forming bonds, this time they only hybridise three of the orbitals rather than all four. They use the 2s electron and two of the 2p electrons, but leave the other 2p electron unchanged. The new orbitals formed are called sp2 hybrids, because they are made by an s orbital and two p orbitals reorganising themselves. sp2 orbitals look rather like sp3 orbitals that you have already come across in the bonding in methane, except that they are shorter and fatter. The three sp2 hybrid orbitals arrange themselves as far apart as possible which is at 120° to each other in a plane. The remaining p orbital is at right angles to them. The two carbon atoms and four hydrogen atoms would look like this before they joined together:
The various atomic orbitals which are pointing towards each other now merge to give molecular orbitals, each containing a bonding pair of electrons. These are sigma bonds just like those formed by end-to-end overlap of atomic orbitals in, say, ethane.
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bonding in ethene - sp2 hybridisation
The p orbitals on each carbon aren't pointing towards each other, and so we'll leave those for a moment. In the diagram, the black dots represent the nuclei of the atoms. Notice that the p orbitals are so close that they are overlapping sideways. This sideways overlap also creates a molecular orbital, but of a different kind. In this one the electrons aren't held on the line between the two nuclei, but above and below the plane of the molecule. A bond formed in this way is called a pi bond.
For clarity, the sigma bonds are shown using lines - each line representing one pair of shared electrons. The various sorts of line show the directions the bonds point in. An ordinary line represents a bond in the plane of the screen (or the paper if you've printed it), a broken line is a bond going back away from you, and a wedge shows a bond coming out towards you.
Note: The really interesting bond in ethene is the pi bond. In almost all cases where you will draw the structure of ethene, the sigma bonds will be shown as lines.
Be clear about what a pi bond is. It is a region of space in which you can find the two electrons which make up the bond. Those two electrons can live anywhere within that space. It would be quite misleading to think of one living in the top and the other in the bottom.
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bonding in ethene - sp2 hybridisation
Beyond A'level: This is another example of the curious behaviour of electrons. How do the electrons get from one half of the pi bond to the other if they are never found in between? It's an unanswerable question if you think of electrons as particles.
Even if your syllabus doesn't expect you to know how a pi bond is formed, it will expect you to know that it exists. The pi bond dominates the chemistry of ethene. It is very vulnerable to attack - a very negative region of space above and below the plane of the molecule. It is also somewhat distant from the control of the nuclei and so is a weaker bond than the sigma bond joining the two carbons.
Important! Check your syllabus! Find out whether you actually need to know how a pi bond is formed. Don't forget to look in the bonding section of your syllabus as well as under ethene. If you don't need to know it, there's no point in learning it! You will, however, need to know that a pi bond exists - that the two bonds between the carbon atoms in ethene aren't both the same. If you haven't got a copy of your syllabus, find out how to download one
All double bonds (whatever atoms they might be joining) will consist of a sigma bond and a pi bond.
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bonding in ethene - sp2 hybridisation
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© Jim Clark 2000
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electronegativity - polar bonds in organic compounds
ELECTRONEGATIVITY
This page deals with electronegativity in an organic chemistry context. If you want a wider view of electronegativity, there is a link at the bottom of the page.
What is electronegativity? Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. The Pauling scale is the most commonly used. Fluorine (the most electronegative element) is given a value of 4.0, and values range down to caesium and francium which are the least electronegative at 0.7. What happens if two atoms of equal electronegativity bond together? The most obvious example of this is the bond between two carbon atoms. Both atoms will attract the bonding pair to exactly the same extent. That means that on average the electron pair will be found half way between the two nuclei, and you could draw a picture of the bond like this:
It is important to realise that this is an average picture. The electrons are actually in a sigma orbital, and are moving constantly within that orbital.
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electronegativity - polar bonds in organic compounds
Help! A sigma orbital is a molecular orbital formed by end-toend overlap between two atomic orbitals. If you aren't happy about this, read the articles on orbitals and the bonding in methane and ethane.
The carbon-fluorine bond Fluorine is much more electronegative than carbon. The actual values on the Pauling scale are carbon fluorine
2.5 4.0
That means that fluorine attracts the bonding pair much more strongly than carbon does. The bond - on average - will look like this:
Why is fluorine more electronegative than carbon? A simple dots-and-crosses diagram of a C-F bond is perfectly adequate to explain it.
The bonding pair is in the second energy level of both carbon and fluorine, so in the absence of any other effect, the distance of the pair from both nuclei would be the same. The electron pair is shielded from the full force of both nuclei by the 1s electrons - again there is nothing to pull it closer to one atom than the
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electronegativity - polar bonds in organic compounds
other. BUT, the fluorine nucleus has 9 protons whereas the carbon nucleus has only 6. Allowing for the shielding effect of the 1s electrons, the bonding pair feels a net pull of about 4+ from the carbon, but about 7+ from the fluorine. It is this extra nuclear charge which pulls the bonding pair (on average) closer to the fluorine than the carbon.
Help! You have to imagine what the bonding pair "sees" if it looks in towards the nucleus. In the carbon case, it sees 6 positive protons, and 2 negative electrons. That means that there will be a net pull from the carbon of about 4+. The shielding wouldn't actually be quite as high as 2-, because the 1s electrons spend some of their time on the far side of the carbon nucleus - and so aren't always between the bonding pair and the nucleus. Incidentally, thinking about electrons looking towards the nucleus may be helpful in picturing what is going on, but avoid using terms like this in exams.
The carbon-chlorine bond The electronegativities are: carbon chlorine
2.5 3.0
The bonding pair of electrons will be dragged towards the chlorine but not as much as in the fluorine case. Chlorine isn't as electronegative as fluorine. Why isn't chlorine as electronegative as fluorine? Chlorine is a bigger atom than fluorine.
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electronegativity - polar bonds in organic compounds
fluorine: 1s22s22px22py22pz1 chlorine: 1s22s22px22py22pz23s23px23py23pz1
Help! If you aren't happy about this, read the article on orbitals. Use the BACK button on your browser to get back to here again.
In the chlorine case, the bonding pair will be shielded by all the 1-level and 2-level electrons. The 17 protons on the nucleus will be shielded by a total of 10 electrons, giving a net pull from the chlorine of about 7+. That is the same as the pull from the fluorine, but with chlorine the bonding pair starts off further away from the nucleus because it is in the 3-level. Since it is further away, it feels the pull from the nucleus less strongly.
Bond polarity and inductive effects Polar bonds Think about the carbon-fluorine bond again. Because the bonding pair is pulled towards the fluorine end of the bond, that end is left rather more negative than it would otherwise be. The carbon end is left rather short of electrons and so becomes slightly positive.
The symbols + and - mean "slightly positive" and "slightly negative". You read + as "delta plus" or "delta positive". We describe a bond having one end slightly positive and the other end slightly negative as being polar. Inductive effects http://www.chemguide.co.uk/basicorg/bonding/eneg.html (4 of 8)30/12/2004 11:02:05
electronegativity - polar bonds in organic compounds
An atom like fluorine which can pull the bonding pair away from the atom it is attached to is said to have a negative inductive effect. Most atoms that you will come across have a negative inductive effect when they are attached to a carbon atom, because they are mostly more electronegative than carbon. You will come across some groups of atoms which have a slight positive inductive effect - they "push" electrons towards the carbon they are attached to, making it slightly negative. Inductive effects are sometimes given symbols: -I (a negative inductive effect) and +I (a positive inductive effect).
Note: You should be aware of terms like "negative inductive effect", but don't get bogged down in them. Provided that you understand what happens when electronegative atoms like fluorine or chlorine are attached to carbon atoms in terms of the polarity of the bonds, that's really all you need for most purposes.
Some important examples of polar bonds Hydrogen bromide (and other hydrogen halides)
Bromine (and the other halogens) are all more electronegative than hydrogen, and so all the hydrogen halides have polar bonds with the hydrogen end slightly positive and the halogen end slightly negative.
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electronegativity - polar bonds in organic compounds
Help! Halogen: a member of group VII of the Periodic Table - fluorine, chlorine, bromine and iodine. Halide: a compound of one of these - e.g. hydrogen chloride, hydrogen bromide, etc.
The polarity of these molecules is important in their reactions with alkenes.
Note: These reactions are explored in the section dealing with the addition of hydrogen halides to alkenes.
The carbon-bromine bond in halogenoalkanes
Note: You may come across halogenoalkanes under the names "haloalkanes" or "alkyl halides".
Bromine is more electronegative than carbon and so the bond is polarised in the way that we have already described with C-F and C-Cl.
The polarity of the carbon-halogen bonds is important in the reactions of the halogenoalkanes.
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electronegativity - polar bonds in organic compounds
Note: This link will take you to the nucleophilic substitution reactions of the halogenoalkanes in which this polarity is important.
The carbon-oxygen double bond An orbital model of the C=O bond in methanal, HCHO, looks like this:
Note: If you aren't sure about this, read the article on bonding in the carbonyl group (C=O).
The very electronegative oxygen atom pulls both bonding pairs towards itself - in the sigma bond and the pi bond. That leaves the oxygen fairly negative and the carbon fairly positive.
Note: You can read about addition reactions or additionelimination reactions of compounds containing carbonoxygen double bonds elswhere on this site.
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electronegativity - polar bonds in organic compounds
Where would you like to go now? To look at electronegativity in a wider context. . . To the organic bonding menu. . . To menu of basic organic chemistry. . . To Main Menu . . .
© Jim Clark 2000
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electrophilic addition - symmetrical alkenes and hydrogen halides
THE REACTION BETWEEN SYMMETRICAL ALKENES AND THE HYDROGEN HALIDES
This page gives you the facts and a simple, uncluttered mechanism for the electrophilic addition reactions between the hydrogen halides and alkenes like ethene and cyclohexene. Hydrogen halides include hydrogen chloride and hydrogen bromide. If you want the mechanisms explained to you in detail, there is a link at the bottom of the page.
Electrophilic addition reactions involving hydrogen bromide The facts Alkenes react with hydrogen bromide in the cold. The double bond breaks and a hydrogen atom ends up attached to one of the carbons and a bromine atom to the other. In the case of ethene, bromoethane is formed.
Note: Be careful when you write the names of the addition products that you change the ene ending in the original alkene (showing the C=C) into an ane ending (showing that it has been replaced by C-C).
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electrophilic addition - symmetrical alkenes and hydrogen halides
With cyclohexene you get bromocyclohexane.
The structures of the cyclohexene and the bromocyclohexane are often simplified:
Note: Each corner in one of these diagrams represents a carbon atom. Each carbon atom has enough hydrogens attached to make the total number of bonds up to 4. In the case of the bromocyclohexane, it isn't necessary to write the new hydrogen into the diagram, but it is helpful to put it there to emphasise that addition has happened.
Be sure that you understand the relationship between these simplified diagrams and the full structures. The mechanisms The reactions are examples of electrophilic addition. With ethene and HBr:
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electrophilic addition - symmetrical alkenes and hydrogen halides
and with cyclohexene:
Electrophilic addition reactions involving the other hydrogen halides The facts Hydrogen chloride and the other hydrogen halides add on in exactly the same way. For example, hydrogen chloride adds to ethene to make chloroethane:
The only difference is in how fast the reactions happen with the different hydrogen halides. The rate of reaction increases as you go from HF to HCl to HBr to HI. HF slowest reaction HCl HBr HI
fastest reaction
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The reason for this is that as the halogen atoms get bigger, the strength of the hydrogen-halogen bond falls. Bond strengths (measured in kilojoules per mole) are: H-F 568 H-Cl 432 H-Br 366 H-I
298
Note: You may find slightly different values depending on which data source you use. It doesn't matter - the differences are minor and the pattern is always the same.
As you have seen in the HBr case, in the first step of the mechanism the hydrogen-halogen bond gets broken. If the bond is weaker, it will break more readily and so the reaction is more likely to happen. The mechanisms The reactions are still examples of electrophilic addition. With ethene and HCl, for example:
This is exactly the same as the mechanism for the reaction between ethene and HBr, except that we've replaced Br by Cl. http://www.chemguide.co.uk/mechanisms/eladd/symhbr.html (4 of 5)30/12/2004 11:02:10
electrophilic addition - symmetrical alkenes and hydrogen halides
All the other mechanisms for symmetrical alkenes and the hydrogen halides would be done in the same way.
Suggestion: Find out if your syllabus mentions a particular hydrogen halide, and learn that mechanism. You can simply swap the halogen atom if a different hydrogen halide comes up in an exam. Haven't got a syllabus? Follow this link to find out how to get one.
Where would you like to go now? Help! Talk me through this mechanism . . . Look at the same reactions involving unsymmetrical alkenes . . . To menu of electrophilic addition reactions. . . To menu of other types of mechanism. . . To Main Menu . . .
© Jim Clark 2000
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Explaining electrophilic addition involving hydrogen halides
EXPLAINING THE REACTION BETWEEN SYMMETRICAL ALKENES AND THE HYDROGEN HALIDES
This page guides you through the mechanism for the electrophilic addition of hydrogen halides such as hydrogen bromide with symmetrical alkenes like ethene or cyclohexene. Unsymmetrical alkenes are covered separately, and you will find a link at the bottom of the page.
Electrophilic addition reactions involving hydrogen bromide Hydrogen bromide is chosen as a typical hydrogen halide. Bromine is more electronegative than hydrogen. That means that the bonding pair of electrons is pulled towards the bromine end of the bond, and so the hydrogen bromide molecule is polar.
Note: If you aren't sure about electronegativity and bond polarity follow this link before you read on. Use the BACK button on your browser to return to this page.
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Explaining electrophilic addition involving hydrogen halides
The reaction of ethene with hydrogen bromide The structure of ethene is shown in the diagram on the right. The pi bond is an orbital above and below the plane of the rest of the molecule, and relatively exposed to things around it. The two electrons in this orbital are highly attractive to anything which is positively charged.
Note: If you aren't sure about this, it would be a good idea to read the introductory page on electrophilic addition before you go on. Use the BACK button on your browser to return to this page.
The slightly positive hydrogen atom in the hydrogen bromide acts as an electrophile, and is strongly attracted to the electrons in the pi bond.
Electrophile: A substance with a strong attraction to a negative region in another substance. Electrophiles are either fully positive ions, or the slightly positive end of a polar molecule.
The electrons from the pi bond move down towards the slightly positive hydrogen atom. In the process, the electrons in the H-Br bond are repelled down until they are entirely on the bromine atom, producing a bromide ion.
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Explaining electrophilic addition involving hydrogen halides
Help! If you aren't sure about the use of curly arrows in mechanisms, you must follow this link before you go on. Use the BACK button on your browser to return to this page.
That leaves you with these two ions at this half-way stage of the reaction:
The ion with a positive charge on the carbon atom is called a carbocation or carbonium ion (an older term). Why is there a positive charge on the carbon atom? The pi bond was originally made up of an electron from each of the carbon atoms. Both of those electrons have been used to make a new bond to the hydrogen. That leaves the right-hand carbon an electron short - hence positively charged. In the second stage of the mechanism, the lone pair of electrons on the bromide ion is strongly attracted to the positive carbon and moves towards it until a bond is formed.
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Explaining electrophilic addition involving hydrogen halides
Note: For clarity only one of the lone pairs around the bromide ion is shown. That's perfectly acceptable, because the other three lone pairs aren't involved in the process - they are pointing in the wrong directions.
The overall mechanism is therefore
The reaction of cyclohexene with hydrogen bromide Cyclohexene is chosen an an example of a fairly commonly used symmetrical alkene. The fact that it is a ring structure doesn't make any difference to the mechanism. The full structure of cyclohexene is
but it is often abbreviated to
In this diagram, there is a carbon atom at each corner, and enough hydrogens attached to each carbon to bring the total number of bonds per carbon atom up to 4.
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Explaining electrophilic addition involving hydrogen halides
The double bond has an easily attacked pi bond exactly as in ethene, and the electrons in that bond are attracted towards the slightly positive hydrogen atom in the HBr. Once again, the pi bond electrons swing to make a bond with the hydrogen, and push the electrons in the H-Br bond fully onto the bromine, making a bromide ion.
Care! Think carefully about which way the pi bond electrons swing. In this case, think of them as being pivotted about the top carbon atom. It is therefore that carbon atom which is joined to the new hydrogen.
In the second stage, one of the lone pairs of electrons on the bromide ion is attracted to the positively charged carbon atom and forms a bond with it.
The overall mechanism is therefore:
Electrophilic addition reactions involving the other hydrogen halides The mechanisms The other hydrogen halides behave in exactly the same way as hydrogen http://www.chemguide.co.uk/mechanisms/eladd/symhbrtt.html (5 of 8)30/12/2004 11:02:15
Explaining electrophilic addition involving hydrogen halides
bromide. For example, compare the reaction between ethene and hydrogen bromide with the one between ethene and hydrogen chloride.
There's no need to learn both mechanisms. As long as you know one of them, all you have to do is swap one halogen atom for another. That's equally true for hydrogen fluoride or hydrogen iodide. The different rates of reaction The rate of reaction increases as you go from HF to HCl to HBr to HI. HF slowest reaction HCl HBr HI
fastest reaction
The reason for this is that as the halogen atoms get bigger, the strength of the hydrogen-halogen bond falls. Bond strengths (measured in kilojoules per mole) are:
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Explaining electrophilic addition involving hydrogen halides
H-F 568 H-Cl 432 H-Br 366 H-I
298
In the first step of these mechanisms, the hydrogen-halogen bond breaks as the electron pair is forced down onto the halogen atom. Breaking bonds needs energy, and if the bond is weaker, it will break more easily - needing less energy. That means that the activation energy for the reactions will fall as you go from hydrogen fluoride to hydrogen iodide. The lower the activation energy, the faster the reaction.
Activation energy: The minimum energy needed before a reaction will occur. In this case it is the energy needed to break the various bonds and make the carbocation and the halide ion.
Beware! A red herring! People sometimes get confused because there is another tempting place to look for the reason why the reaction rates are different between the various hydrogen halides. The halogens have different electronegativities - with fluorine being the most electronegative and iodine the least. That means that the hydrogen in HF will have the greatest positive charge and so will be attracted most strongly to the pi bond. It would be tempting to think that that would produce the fastest reaction
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Explaining electrophilic addition involving hydrogen halides
- but not so! Although the HF may well be attracted most strongly, attraction alone isn't enough. If anything is to happen, bonds have to be broken - and here HF is at a disadvantage, because the bond is very strong. The lesson from all this When you are trying to find reasons for differing rates of reactions, always look first at differences in bond strengths. Electronegativity differences may be interesting, but rarely give you the answer you want!
Where would you like to go now? Look at the same reactions involving unsymmetrical alkenes . . . To menu of electrophilic addition reactions. . . To menu of other types of mechanism. . . To Main Menu . . .
© Jim Clark 2000
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The use of curly arrows to show electron movements in reaction mechanisms
USING CURLY ARROWS IN REACTION MECHANISMS
This page explains the use of curly arrows to show the movement both of electron pairs and of single electrons during organic reaction mechanisms. You can jump straight to the movement of single electrons further down this page if that is all you are interested in for the moment (for example, if you are currently working on free radical reactions).
Using curly arrows to show the movement of electron pairs Curly arrows (and that's exactly what they are called!) are used in mechanisms to show the various electron pairs moving around. You mustn't use them for any other purpose. ●
●
The arrow tail is where the electron pair starts from. That's always fairly obvious, but you must show the electron pair either as a bond or, if it is a lone pair, as a pair of dots. Remember that a lone pair is a pair of electrons at the bonding level which isn't currently being used to join on to anything else. The arrow head is where you want the electron pair to end up.
For example, in the reaction between ethene and hydrogen bromide, one of the two bonds between the two carbon atoms breaks. That bond is simply a pair of electrons. Those electrons move to form a new bond with the hydrogen from the HBr. At the same time the pair of electrons in the hydrogen-bromine bond moves down on to the bromine atom.
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The use of curly arrows to show electron movements in reaction mechanisms
There's no need to draw the pairs of electrons in the bonds as two dots. Drawing the bond as a line is enough, but you could put two dots in as well if you wanted to. Notice that the arrow head points between the C and H because that's where the electron pair ends up. Notice also that the electron movement between the H and Br is shown as a curly arrow even though the electron pair moves straight down. You have to show electron pair movements as curly arrows - not as straight ones. The second stage of this reaction nicely illustrates how you use a curly arrow if a lone pair of electrons is involved. The first stage leaves you with a positive charge on the right hand carbon atom and a negative bromide ion. You can think of the electrons shown on the bromide ion as being the ones which originally made up the hydrogen-bromine bond.
Note: There are another three lone pairs around the outside of the bromide ion - making four in all. These aren't normally shown because they don't actually do anything new and interesting! However, it is essential that you show the lone pair you are interested in as a pair of dots. If you don't, you risk losing marks in an exam.
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The use of curly arrows to show electron movements in reaction mechanisms
The lone pair on the bromide ion moves to form a new bond between the bromine and the right hand carbon atom. That movement is again shown by a curly arrow. Notice again, that the curly arrow points between the carbon and the bromine because that's where the electron pair ends up. That leaves you with the product of this reaction, bromoethane:
Note: You can read a full description of this mechanism together with other similar reactions of ethene and the other alkenes by following this link.
Using curly arrows to show the movement of single electrons The most common use of "curly arrows" is to show the movement of pairs of electrons. You can also use similar arrows to show the movement of single electrons - except that the heads of these arrows only have a single line rather than two lines. shows the movement of an electron pair
shows the movement of a single electron The first stage of the polymerisation of ethene, for example, could be shown as:
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The use of curly arrows to show electron movements in reaction mechanisms
You should draw the dots showing the interesting electrons. The half arrows show where they go. This is very much a "belt-and-braces" job, and the arrows don't add much. Whether you choose to use these half arrows to show the movement of a single electron should be governed by what your syllabus says. If your syllabus encourages the use of these arrows, then it makes sense to use them. If not - if the syllabus says that they "may" be used, or just ignores them altogether - then they are as well avoided. There is some danger of confusing them with the arrows showing electron pair movements, which you will use all the time. If, by mistake, you use an ordinary full arrow to show the movement of a single electron you run the risk of losing marks.
Help! You must have a copy of your syllabus! If you haven't got a copy, find out how to get a syllabus by following this link.
Where would you like to go now? To the organic conventions menu. . . To menu of basic organic chemistry. . . To Main Menu . . .
© Jim Clark 2000
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electrophilic addition - unsymmetrical alkenes and hydrogen halides
THE REACTION BETWEEN UNSYMMETRICAL ALKENES AND THE HYDROGEN HALIDES
This page gives you the facts and a simple, uncluttered mechanism for the electrophilic addition reactions between the hydrogen halides and alkenes like propene. Hydrogen halides include hydrogen chloride and hydrogen bromide. If you want the mechanisms explained to you in detail, there is a link at the bottom of the page. An unsymmetrical alkene is one like propene in which the groups or atoms attached to either end of the carbon-carbon double bond are different. For example, in propene there are a hydrogen and a methyl group at one end, but two hydrogen atoms at the other end of the double bond. But-1ene is another unsymmetrical alkene.
Electrophilic addition reactions involving hydrogen bromide The facts As with all alkenes, unsymmetrical alkenes like propene react with hydrogen bromide in the cold. The double bond breaks and a hydrogen atom ends up attached to one of the carbons and a bromine atom to the other. In the case of propene, 2-bromopropane is formed.
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electrophilic addition - unsymmetrical alkenes and hydrogen halides
This would normally be written in a more condensed form as
The product is 2-bromopropane.
Note: There is another possible reaction between unsymmetrical alkenes and hydrogen bromide (but not the other hydrogen halides) unless the hydrogen bromide and alkene are absolutely pure. A different mechanism happens (a free radical chain reaction - not on UK A' level syllabuses) which leads to the hydrogen and bromine adding the opposite way round. For A' level purposes, you don't need to worry about that. However, if you are interested, you will find the free radical addition mechanism by following this link. Use the BACK button on your browser to return to this page later.
This is in line with Markovnikov's Rule which says: When a compound HX is added to an unsymmetrical alkene, the hydrogen becomes attached to the carbon with the most hydrogens attached to it already. In this case, the hydrogen becomes attached to the CH2 group, because the CH2 group has more hydrogens than the CH group. Notice that only the hydrogens directly attached to the carbon atoms at either end of the double bond count. The ones in the CH3 group are http://www.chemguide.co.uk/mechanisms/eladd/unsymhbr.html (2 of 5)30/12/2004 11:02:26
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totally irrelevant.
Warning! Markovnikov's Rule is a useful guide for you to work out which way round to add something across a double bond, but it isn't the reason why things add that way. As a general principle, don't quote Markovnikov's Rule in an exam unless you are specifically asked for it.
The mechanism This is an example of electrophilic addition.
The addition is this way around because the intermediate carbocation (previously called a carbonium ion) formed is secondary. This is more stable (and so is easier to form) than the primary carbocation which would be produced if the hydrogen became attached to the centre carbon atom and the bromine to the end one.
Electrophilic addition reactions involving the other hydrogen halides The facts Hydrogen fluoride, hydrogen chloride and hydrogen iodide all add on in exactly the same way as hydrogen bromide. The only differences lie in the rates of reaction:
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electrophilic addition - unsymmetrical alkenes and hydrogen halides
HF slowest reaction HCl HBr HI
fastest reaction
This is because the hydrogen-halogen bond gets weaker as the halogen atom gets bigger. If the bond is weaker, it breaks more easily and so the reaction is faster. If the halogen is given the symbol X, the equation for the reaction with propene is:
Notice that the product is still in line with Markovnikov's Rule. The mechanism These are still examples of electrophilic addition. Again using X to stand for any halogen:
Again, the intermediate carbocation formed is secondary. This is more stable than the primary carbocation ion which would be formed if the hydrogen attached to the centre carbon atom and the X to the end one. If http://www.chemguide.co.uk/mechanisms/eladd/unsymhbr.html (4 of 5)30/12/2004 11:02:26
electrophilic addition - unsymmetrical alkenes and hydrogen halides
it is more stable it will be easier to make.
Where would you like to go now? Help! Talk me through this mechanism . . . To menu of electrophilic addition reactions. . . To menu of other types of mechanism. . . To Main Menu . . .
© Jim Clark 2000
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hydrogen bromide and alkenes - the peroxide effect
HYDROGEN BROMIDE AND ALKENES: THE PEROXIDE EFFECT
This page gives you the facts and simple uncluttered mechanisms for the free radical addition of hydrogen bromide to alkenes - often known as the "peroxide effect". If you want the mechanisms explained to you in more detail, there is a link at the bottom of the page.
Addition to symmetrical alkenes A symmetrical alkene is one like ethene where the groups at both ends of the carbon-carbon double bond are the same. The facts The reaction happens at room temperature in the presence of organic peroxides or some oxygen from the air. Alkenes react very slowly with oxygen to produce traces of organic peroxides - so the two possible conditions are equivalent to each other. The reaction is a simple addition of the hydrogen bromide. For example, with ethene:
With a symmetrical alkene you get exactly the same product in the absence of the organic peroxides or oxygen - but the mechanism is different.
The mechanism Hydrogen halides (hydrogen chloride, hydrogen bromide and the rest) usually react with alkenes using an electrophilic addition mechanism. However, in the presence of organic peroxides, hydrogen bromide adds http://www.chemguide.co.uk/mechanisms/freerad/alkenehbr.html (1 of 6)30/12/2004 11:02:33
hydrogen bromide and alkenes - the peroxide effect
by a different mechanism.
Note: If you are interested, you will find the electrophilic addition mechanism for the addition of hydrogen bromide and other hydrogen halides to alkenes if you follow this link. You may need to explore several pages in this section. Use the BACK button (or the HISTORY file or GO menu) on your browser to return to this page.
With the organic peroxides present you get a free radical chain reaction. Chain initiation The chain is initiated by free radicals produced by an oxygen-oxygen bond in the organic peroxide breaking.
These free radicals extract a hydrogen atom from a hydrogen bromide molecule to produce bromine radicals.
Chain propagation A bromine radical joins to the ethene using one of the electrons in the pi bond. That creates a new radical with the single electron on the other carbon atom.
That radical reacts with another HBr molecule to produce bromoethane and another bromine radical to continue the process.
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hydrogen bromide and alkenes - the peroxide effect
etc Chain termination Eventually two free radicals hit each other and produce a molecule of some sort. The process stops here because no new free radicals are formed.
Addition to unsymmetrical alkenes An unsymmetrical alkene is one like propene where the groups at either end of the carbon-carbon double bond are different. The facts The reaction happens under the same conditions as with a symmetrical alkene, but there is a complication because the hydrogen and the bromine can add in two different ways. Which way they add depends on whether there are organic peroxides (or oxygen) present or not.
Normally, when a molecule HX adds to a carbon-carbon double bond, the hydrogen becomes attached to the carbon with the more hydrogens on already. This is known as Markovnikov's Rule. Because the HBr adds on the "wrong way around " in the presence of http://www.chemguide.co.uk/mechanisms/freerad/alkenehbr.html (3 of 6)30/12/2004 11:02:33
hydrogen bromide and alkenes - the peroxide effect
organic peroxides, this is often known as the peroxide effect or antiMarkovnikov addition. In the absence of peroxides, hydrogen bromide adds to propene via an electrophilic addition mechanism. That gives the product predicted by Markovnikov's Rule.
The free radical mechanism Chain initiation This is exactly the same as in the ethene case above.
Chain propagation When the bromine radical joins to the propene, it attaches so that a secondary radical is formed. This is more stable (and so easier to form) than the primary radical which would be formed if it attached to the other carbon atom.
That radical reacts with another HBr molecule to produce 1bromopropane and another bromine radical to continue the process.
etc Chain termination Eventually two free radicals hit each other and produce a molecule of some sort. The process stops here because no new free radicals are formed.
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hydrogen bromide and alkenes - the peroxide effect
Why don't the other hydrogen halides behave in the same way? The reason that hydrogen bromide adds in an anti-Markovnikov fashion in the presence of organic peroxides is simply a question of reaction rates. The free radical mechanism is much faster than the alternative electrophilic addition mechanism. Both mechanisms happen, but most of the product is the one from the free radical mechanism because that is working faster. With the other hydrogen halides, the opposite is true. Hydrogen fluoride The hydrogen-fluorine bond is so strong that fluorine radicals aren't formed in the initiation step. Hydrogen chloride With hydrogen chloride, the second half of the propagation stage is very slow. If you do a bond enthalpy sum, you will find that the following reaction is endothermic.
This is due to the relatively high hydrogen-chlorine bond strength. Hydrogen iodide In this case, the first step of the propagation stage turns out to be endothermic and this slows the reaction down. Not enough energy is released when the weak carbon-iodine bond is formed.
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hydrogen bromide and alkenes - the peroxide effect
exothermic.
Where would you like to go now? Help! Talk me through these mechanisms . . . To menu of free radical reactions. . . To menu of other types of mechanism. . . To Main Menu . . .
© Jim Clark 2003
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explaining the peroxide effect in the addition of hydrogen bromide to alkenes
EXPLAINING THE "PEROXIDE EFFECT" IN THE REACTION BETWEEN HYDROGEN BROMIDE AND ALKENES
This page guides you through the mechanism for the free radical addition of hydrogen bromide to alkenes - often known as the "peroxide effect".
Note: If you just want the facts and mechanism with a minimum of discussion you will find them by following this link.
Free radical addition to a carbon-carbon double bond If you have read the introductory page (see above), you will know that hydrogen bromide adds to the carbon-carbon double bond in alkenes via a free radical mechanism in the presence of organic peroxides or oxygen from the air. Oxygen reacts slowly with alkenes to produce small amounts of organic peroxides, so we don't need to look at that as a separate case. We'll start by looking at the general case without worrying about what is attached at either end of the carbon-carbon double bond.
The function of the organic peroxides - chain initiation Organic peroxides are compounds containing an oxygen-oxygen single bond, and are commonly given a general formula R-O-O-R. The "R" groups can be quite complicated and aren't necessarily just simple alkyl groups. http://www.chemguide.co.uk/mechanisms/freerad/alkenehbrtt.html (1 of 8)30/12/2004 11:02:41
explaining the peroxide effect in the addition of hydrogen bromide to alkenes
The oxygen-oxygen bond is quite weak, and breaks easily so that each oxygen gets a single electron. Free radicals are formed.
If these free radicals collide with a hydrogen bromide molecule, a hydrogen atom is transferred, breaking the hydrogen-bromine bond to produce bromine radicals.
Chain propagation In any alkene (like ethene, for example), the two pairs of electrons which make up the double bond aren't the same. One pair is held securely on the line between the two carbon nuclei in a bond called a sigma bond. The other pair is more loosely held in an orbital above and below the plane of the molecule known as a pi bond.
Note: It would be helpful - but not essential - if you read about the structure of ethene before you went on. If the diagram above is unfamiliar to you, then you certainly ought to read this background material. Use the BACK button on your browser to return to this page.
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explaining the peroxide effect in the addition of hydrogen bromide to alkenes
Imagine what happens if a free radical approaches the pi bond in an alkene. Once again, we'll draw it as if it is ethene - but it would apply to any case. The bromine radical uses one of the electrons in the pi bond to help to form a new bond between itself and the left hand carbon atom. The other electron returns to the right hand carbon.
Note: Don't worry that we've gone back to a simpler diagram. It is perfectly adequate for this discussion.
The sigma bond between the carbon atoms isn't affected by any of this. Now this new free radical reacts with a hydrogen bromide molecule. It takes a hydrogen atom from it, leaving another bromine radical.
The bromine radical can then react with another carbon-carbon double bond, which eventually produces a new bromine radical - and so on, and so on . . . There is a chain reaction.
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explaining the peroxide effect in the addition of hydrogen bromide to alkenes
Chain termination The chain will be broken when any two radicals happen to hit each other and form a new bond using both of the single electrons. Removing a free radical from the system without producing a new one immediately stops that particular chain.
What happens if the alkene is unsymmetrical? Unsymmetrical alkenes? An unsymmetrical alkene is one like propene, CH3CH=CH2. At one end of the double bond there is a CH3 group and a hydrogen atom. At the other end there are two hydrogen atoms. A question of orientation The problem with these unsymmetrical alkenes is that you could get two different products depending on which end of the bond the hydrogen and the bromine add. In fact, under free radical conditions, most of the product is 1bromopropane.
This happens because when the bromine radical attacks the pi bond, it joins to the carbon atom at the CH2 end of the double bond rather than the CH end.
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explaining the peroxide effect in the addition of hydrogen bromide to alkenes
Why does the bromine add this way? You might think that there would be an equal chance of it attaching to either end, but where it attaches is controlled by the stability of the free radical formed. The more stable radical will be formed more quickly. Think of this in terms of activation energy.
The activation energy will be lower for the reaction where the bromine attaches to the end carbon because the radical produced is more energetically stable.
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explaining the peroxide effect in the addition of hydrogen bromide to alkenes
The stability of various sorts of radicals What matters is the number of carbon atoms attached to the carbon with the single electron. Looking at the simplest possibilities:
Tertiary radicals are more stable than secondary ones, and secondary radicals are more stable than primary. In the case of a bromine radical attacking the double bond in propene, it forms a secondary radical rather than a primary one because it is more stable.
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explaining the peroxide effect in the addition of hydrogen bromide to alkenes
Note: The order of stability of the various tertiary, secondary and primary free radicals exactly reflects the order of stability of carbocations (carbonium ions). If you are interested in following this link, use the BACK button on your browser to return to this page.
The rest of the reaction Once the bromine has attached to the carbon to form the secondary radical, there is nothing different about the rest of the reaction. The new radical takes a hydrogen from a hydrogen bromide molecule. This produces the 1-bromopropane and a bromine radical.
The bromine radical now goes into another cycle exactly as before to continue the chain reaction.
Where would you like to go now? To menu of free radical reactions. . . To menu of other types of mechanism. . . To Main Menu . . .
© Jim Clark 2003
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explaining the peroxide effect in the addition of hydrogen bromide to alkenes
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carbocations (or carbonium ions)
CARBOCATIONS (or CARBONIUM IONS) All carbocations (previously known as carbonium ions) carry a positive charge on a carbon atom. The name tells you that - a cation is a positive ion, and the "carbo" bit refers to a carbon atom. However there are important differences in the structures of various types of carbocations.
The different kinds of carbocations Primary carbocations In a primary (1°) carbocation, the carbon which carries the positive charge is only attached to one other alkyl group.
Help! An alkyl group is a group such as methyl, CH3, or ethyl, CH3CH2. These are groups containing chains of carbon atoms which may be branched. Alkyl groups are given the general symbol R.
Some examples of primary carbocations include:
Notice that it doesn't matter how complicated the attached alkyl group is. All you are doing is counting the number of bonds from the positive carbon to other carbon atoms. In all the above cases there is only one such link.
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carbocations (or carbonium ions)
Using the symbol R for an alkyl group, a primary carbocation would be written as in the box.
Secondary carbocations In a secondary (2°) carbocation, the carbon with the positive charge is attached to two other alkyl groups, which may be the same or different. Examples:
A secondary carbocation has the general formula shown in the box. R and R' represent alkyl groups which may be the same or different.
Tertiary carbocations In a tertiary (3°) carbocation, the positive carbon atom is attached to three alkyl groups, which may be any combination of same or different.
A tertiary carbocation has the general formula shown in the box. R, R' and R" are alkyl groups and may be the same or different.
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carbocations (or carbonium ions)
The "electron pushing effect" of alkyl groups You are probably familiar with the idea that bromine is more electronegative than hydrogen, so that in a H-Br bond the electrons are held closer to the bromine than the hydrogen. A bromine atom attached to a carbon atom would have precisely the same effect - the electrons being pulled towards the bromine end of the bond. The bromine has a negative inductive effect.
Help! If you aren't familiar with all of this, follow this link to read about electronegativity and bond polarity before you go any further. Use the BACK button on your browser to return to this page.
Alkyl groups do precisely the opposite and, rather than draw electrons towards themselves, tend to "push" electrons away.
Note: The term "electron pushing" is only to help remember what happens. The alkyl group doesn't literally "push" the electrons away - the other end of the bond attracts them more strongly.
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carbocations (or carbonium ions)
This means that the alkyl group becomes slightly positive ( +) and the carbon they are attached to becomes slightly negative ( -). The alkyl group has a positive inductive effect. This is sometimes shown as, for example:
The arrow shows the electrons being "pushed" away from the CH3 group. The plus sign on the left-hand end of it shows that the CH3 group is becoming positive. The symbols + and - simply reinforce that idea. The importance of spreading charge around in making ions stable The general rule-of-thumb is that if a charge is very localised (all concentrated on one atom) the ion is much less stable than if the charge is spread out over several atoms. Applying that to carbocations of various sorts . . .
You will see that the electron pushing effect of the CH3 group is placing more and more negative charge on the positive carbon as you go from primary to secondary to tertiary carbocations. The effect of this, of course, is to cut down that positive charge. At the same time, the region around the various CH3 groups is becoming somewhat positive. The net effect, then, is that the positive charge is being spread out over more and more atoms as you go from primary to secondary to tertiary ions. The more you can spread the charge around, the more stable the ion becomes.
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carbocations (or carbonium ions)
Order of stability of carbocations primary < secondary < tertiary
Note: The symbol "
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