Unit 1 Notes For Edexcel Chemistry AS Level

 

 AS Chemistry Edexcel Unit 1 Core Principles

 

Unit 1.3 Atomic structure structure & the the periodic table (p52 –  77  77 in text book) 1.  Relative atomic mass The relative atomic mass of an element is the average mass of its isotopes compared with the mass of an atom of carbon-12. 2.  Relative formula mass For a compound, the relative formula mass is the sum of the relative atomic masses of all the atoms in the chemical formula. 3.  Measuring the mass of an atom To find the relative atomic mass of an atom, you need to measure its mass and compare it with the mass of an atom of carbon-12. carbon-12. This can’t can’t be done by weighing. a. 

The mass spectrometer

This is the instrument used to find out the mass of a sample. (Refer to fig 1.3.1 p52 for diagram) There are several different stages to the mass spectrometer; i. 

ii.

Electron gun – gun –  this produces

ii. 

 Acceleration –  Acceleration  – achieved  achieved by

iii. 

 Velocity selector selector –  – to  to ensure

iv. 

vi.

 Vaporisation –  Vaporisation  – this  this is necessary to

due to

Deflection – Deflection  – achieved  achieved by

Detection – Detection  – shows  shows the abundance of ions with each different mass:charge ratio (m/z value)

 

b. 

Finding the relative atomic mass from a mass spectrum

(Refer to fig 1.3.2 p53)  Ar of sample = (m/z x % abundance) abundance) + (m/z (m/z x % abundance) abundance) + …  …  (Refer to green box p54) c. 

Uses of mass spectrometry

(Refer to HSW boxes p55-59) i.  ii.  iii. 

Radioactive dating Drug testing Space exploration

4.  The arrangement of electrons in atoms (Refer to HSW boxes p60-61) Over time, ideas about atomic structure changed as technology improved. a. 

Thomson’s plum pudding model

(Refer to fig 1.3.13) He knew that atoms were neutral but that they contained electrons with a negative charge. He suggested the positive positive ‘pudding’ ‘pudding’ had negative electrons electrons embedded in the structure. He also established the mass of an electron e lectron as approximately 1/2000 of a hydrogen atom. Further work proved that the positive protons have a larger mass than that of electrons. b. 

Rutherford’s atom with a nucleus  nucleus  

(Refer to fig 1.3.15 p61) Rutherford beamed beamed alpha particles particles through gold foil. While most of the the particles were detected detected on the other side, side, some were deflected. deflected. From this he he deduced that there was a central ce ntral positive positive core (the nucleus) surrounded by electrons. He was able to use his measurements measurements to calcu calculate late the diameter of the nucleus and hence the diameter of an atom.

 

c. 

Bohr’s electron shells  shells 

If gas is heated, or electrically charged, it gives out light which can then be split through a prism or diffraction grating to form a spectrum known as a line spectrum or emission spectrum. This is always identical for any any given atom. (Refer to fig 1.3.16 p61) Bohr came up with the idea that electrons were in shells around the nucleus. Moving out from the nucleus successive shells got closer together, in the same way that lines from the emission spectra for hydrogen did. Bohr’s new model matched exactly the line spectrum for hydrogen. Unfortunately, this this only works for hydrogen. hydrogen. Heisenberg & Schrödinger Schrödinger continued this work, beginning quantum mechanics. 5.  Energy levels & electron shells The absorption and emission of light by an atom can be explained e xplained by electron movements between different fixed energy levels i.e. electron shells. Light is a form of energy. When atoms receive receive this light energy absorption absorption takes place and the electrons become ‘excited’ become ‘excited’ and are ‘promoted’ ‘promoted’ to a higher energy level (shell). Conversely, emis emission sion occurs when electrons electrons fall back to a lower energy level (shell). (Refer to fig 1.3.18 p62) Each energy level (shell) is given given a principal quantum quantum number, n. So the ground state for hydrogen is n = 1, since hydrogen has its electron on the first shell (energy level). level). Each successive energy level rises in num number ber by one each time (2,3,4,5….to infinity).  infinity).  6.  Ionisation energies Ionisation is the complete removal of an electron from an atom, which is an endothermic process since energy must be used to overcome the forces of attraction between the electron and the nucleus. Comparing the ionisation energies of different atoms enables a picture to be built up of the electronic e lectronic structure for each atom.  An atom in its ground state state is at its lowest energy level. The energy required required to remove the first electron is called the first ionisation energy, the energy required to remove the second electron is called the second ionisation ionis ation energy and so on. To calculate the total energy energy of removing removing two electrons you must must add the first and second ionisation energies together.

 

7.  Subshells Each electron shell may contain several different subshells. These are described by letters: letters: s, p, d, f, g. The following subshells subshells are available in each shell: Shell

Subshell(s)

1 2 3 4

1s 2s, 2p 3s, 3p, 3d 4s, 4p, 4d, 4f

Within a shell, the subshells have different energy levels, with electrons in the lowest energy subshells being closest to the nucleus: s (lowest energy) < p < d Each type of subshell contains one or more orbitals. First shell (n=1)

Second shell (n=2)

Third shell (n=3)

Fourth shell (n=4)

Subshells

s

sp

spd

spdf

Number of orbitals

1

13

135

1357

8.  Electron spin (Refer to fig 1.3.20 p64) Electrons in atoms behave behave like tiny magnets. magnets. A moving charge can create create a magnetic field. Electrons can spin spin either either clockwise clockwise or anticlockwise. anticlockwise. Two electrons in the same orbital cannon have the the same spin. This means that that each orbital can have a maximum of 2 electrons, having opposite spin. Subshell s p d f

No of orbitals 1 3 5 7

Max No of electrons in subshell 2 6 10 14

 

9.  Electron configurations (filling the orbitals) Each orbital is filled in order, order, lowest en energy ergy orbital first. This means a shell is not always completely filled before electrons start filling the next shell e.g. a 4s orbital will fill before before a 3d orbital. The order of filling is 1s, 2s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p…  6p…  (Refer to fig 1.3.21 p65) Hund’s rule states that electrons in the 2p subshells are placed in different orbitals, rather than placing them in the same orbital with opposite spin. They ‘spread out’ to maximise maximise the number number of unpaired electrons. Once each orbital has had a single unpaired electron placed in it, then the electrons begin to pair up until until the subshell is comp completely letely filled. (This is true of all subshells with multiple orbitals.) (Refer to green box p65 ) There are a few exceptions to this in order to reach a more stable energy state e.g. Copper

Chromium

10. 

Electron configurations for ions

For an ion, you simply add or subtract the number of electrons to form the ion required. 11. 

Shorthand Shorthan d notations

The previous noble gas can be used to avoid writing out long electron configurations. (Refer to HSW box p66)

 

12. 

Electron density maps

Modern scientists scientists regard electrons electrons as waves rather rather than particles. An electron behaves as though it is spread out in an electron cloud surrounding the nucleus. Quantum mechanics mechanics calculations provide provide electron density m maps aps which plot the most likely likely positions of the electrons. An orbital is the the region where the probability of finding an electron is greatest. (Refer to fig 1.3.22 p66) From these maps, scientists have determined the shapes and orientations of orbitals within subshells. (Refer to fig 1.3.23 p67) 13. 

Structure of the periodic table

Elements are arranged in order order of atomic number. number. Each column is called a group, while each row is called a period.  All elements in the the same period period have the the same number number of electron shells i.e. the same quantum number (e.g. n=3 for period 3). The elements in each group or period show similar characteristics in either chemical or physical behaviour. 14. 

The development of the periodic table

The periodic table took on several different forms before the current model was settled upon. (Refer to HSW boxes p70-71) a. 

Early ideas about atomic mass

1799 – Joseph 1799 –  Joseph Proust showed that mass proportions for each element remained constant in compounds 1800 – 1800  – John  John Dalton made the initial proposal that elements were made up of atoms. Unfortunately, Unfortunately, Dalton wasn’t very good at working out the atomic masses 1828 – 1828  – Jons  Jons Jacob Berzelius published an accurate list of atomic masses, but thanks to Dalton’s earlier mistakes, no one took any interest  interest  b. 

Doberiener’s triads  triads 

1829 – Johann 1829 –  Johann Doberiener showed that many of the then known elements could be arranged in groups groups of three i.e. triads. This was the first time time elements had been grouped together

 

1862 – Alexandre-Emile 1862 –  Alexandre-Emile Beguyer de Chancourtois showed similarities between every eighth element. element. However, the dia diagram gram wasn’t published published with his work causing it to be ignored. c. 

Newlands’  law  law of octaves

1863 - John Newlands a arranged rranged ele elements ments in groups of eight by atomic mass. mass. His work was dismissed since it contained some major flaws. He’d assumed all elements had been discovered so didn’t account for any gaps.  gaps.   d. 

Mendeleev’s periodic table  table 

1869 – Dmitri 1869 –  Dmitri Mendeleev and Julius Meyer published clear presentations. Meyer plotted physical properties against atomic mass to produce curves demonstrating demonstratin g periodic relationships. relationships. However, Mendeleev Mendeleev published first producing a table with gaps in it for elements yet to be discovered. He predicted the properties of the missing elements and when more were discovered their properties properties matched matched his predictions. Consequently, Consequently, Mendeleev is credited with producing the modern periodic table. 15. 

The blocks of the periodic table

The periodic table can be divided into blocks determined by the electron orbitals which are being filled. a. 

The s-block elements

This is made up of groups 1 and 2 since their outer electrons e lectrons are in an s orbital. These elements elements are reactive metals since the s electrons electrons are easily lost. They have low melting & boiling points, points, and lower densities densities than other metals but can still conduct electricity. Hydrogen and helium are also s-block elements using these criteria. However, both are non-metals and so scientists treat them as a separate group. b. 

The d-block elements

These are called transition metals, although some of the outer d-block d -block elements don’t share the same characteristics, e.g. zinc. They are much less reactive since the d orbitals are being filled after the outermost s orbital. They tend to follow the general properties of metals, e.g. ductile, malleable, sonorous, conductors conductors etc. Some also make very good good catalysts. catalysts. c. 

The f-block elements

 

The top row (lanthanides) (lanthanides) are all similar similar metals. The bottom row (actinides) (actinides) are radioactive. Some have been synthesised and and are very unstable. unstable. d. 

The p-block elements

This block contains all the non-metals (except hydrogen & helium), along with the metalloids (semi-metals) and some metals. The p-block p-block metals are relatively unreactive and doesn’t exhibit exhibit strong metal character. The metalloids occur in a diagonal diagonal block. They generally behave behave as nonmetals but can conduct electricity e.g. silicon. The non-metals form covalent bonds with each other and ionic bonds with metals. Most non-metals non-metals exist as as small molecules. molecules. The noble gases are extremely extremely unreactive. However, some some compounds have have been formed. (Refer to HSW box p73) 16. Trends Trends in the periodic table Physical and chemical trends occur both across a period and down a group within the periodic table. (Any specific examples are based on period 3) a.   Atomic radius In period 3, all the outer electrons electrons are in the third en energy ergy level: n=3. Going across the period, a proton proton is added to to the nucleus of each element. element. This increases the nuclear charge and therefore increases the attraction between the nucleus and the the electrons. The electrons are pulled pulled closer to the nucleus, nucleus, thereby decreasing the size of the atomic radius. Going down a group

 

b.  Ionic radius For the metal ions – ions – the  the number of electrons decreases, while the proton number increases across across the period. period. This means the the electrons are pulled pulled much closer to the nucleus, decreasing the ionic radius. For the non-metal ions – ions – the  the number of electrons increases as the number of protons increases. The nucleus is less ab able le to pull the electrons towards towards it, so the ionic radius increases. Going down a group

c.  Periodic trends in ionisation energy (Refer to fig 1.3.33 p76) The general trend shows an increase increase across the period. period. As you move move across the period the nucleus of the atoms contains more protons and therefore a higher charge density. density. This results in the the electrons being a attracted ttracted further to the nucleus across the group making the ionisation energy increase. d.  Melting & boiling points (Refer to fig 1.3.34 p77) The general trend shows a decrease across the period but there is a spike for silicon. This can be explained by the the bonding bonding of the elements. Na

 Al

Increases because the atoms form ions which are surrounded by a

sea of delocalised electrons. electrons. Across the the period the charge density of the ions increases causing an increase in electrostatic attractions within the lattice structure. Silicon forms a macromolecular structure structure with strong covalent bonds that need to be broken to change state. The other elements form simple covalent molecules. molecules. The stronger the intermolecular forces (generally stronger for larger molecules) the higher the melting & boiling points.

 

Unit 1.4 Bonding (p78 –  93  93 in text book ) 1. 

What is a chemical bond?

 A bond is the the force holding together atoms atoms together. The physical and and chemical properties of a molecule depend on the type of bond holding together the atoms. 2. 

Ionic bonding

These are formed by metals and non-metals bonding together via electrostatic forces forces of attraction. Consequently Consequently they are also also known as electrovalent bonds. bonds. The ions then become become arranged in a giant lattice structure. Within the lattice lattice the attractive forces between between oppositely oppositely charged ions are maximised whilst the repulsive forces between like charges are minimised. The forces within the lattice lattice therefore act equally equally in all directions. (Refer to fig 1.4.2 p78) Ions are formed when electrons electrons are either lost or gained. gained. The movement of electrons seeks to obey the octet rule;  ‘When elements elements react, they tend to do so in a way that results in an outer shell containing eight electrons.’  ele ctrons.’   Example Na+  1s2 2s2 2p6  This then becomes isoelectronic with with Ne, a noble noble gas (Ne 1s 2 2s2 2p6) i.e. they share the same number and arrangement of electrons. elec trons. Further example Cl-   Ar So, Na loses an electron elec tron to obey the octet rule, whilst Cl gains an electron to do so. When Na and Cl react re act together to form NaCl, the ionic substance produced has different properties to the constituent elements due to the lattice structure.

 

3. 

Dot and cross diagrams

These show the transfer of electrons between atoms in order to form ionic compounds. Only the outer electrons electrons are shown in order order to simplify the diagrams (inner electrons do not get directly involved in tthe he bonding process). The dots and crosses distinguish between electrons of the cations and anions. (Cations are positive ions, while anions are negative ions.) (Refer to fig 1.4.3 p79) 4. 

Trends in ionic radii

The ionic radius is the radius radius of an ion in a crystal. The radius of a cation is smaller than the element’s atomic radius, whilst the radius of an anion is larger than the element’s e lement’s atomic radius. (Refer to fig 1.4.4 p80) 5. 

Types of lattice structure

X-ray diffraction can produce electron density maps of llattice attice arrangements. arrangements. These show that the exact arrangement of ions in a lattice can vary depending on the relative sizes of the different ions present. Sodium chloride exhibits a face-centred cubic structure. Each ion has 6 neighbouring ions, ions, so has a coordination n number umber of 6. This is a common ionic structure. (Refer to fig 1.4.5a p81) Caesium chloride exhibits a body-centred cubic structure since caesium is a larger cation than than sodium. More chloride ions can fit around each cation, cation, so the coordination number in this instance is 8. (Refer to fig 1.4.5b p81) 6. 

Evidence for the existence of ions

(Refer to HSW boxes p82-83) Physical properties i. 

Electrical conductivity

Ionic solids do not not condu conduct ct electricity. This is because because they are in a rigid lattice structure.

 

However, ionic compounds will conduct electricity when either molten or in solution since the ions are free to move. (Refer to fig 1.4.6 p82) ii. 

Strength

Unlike metals, ionic compounds compounds are not not malleable or ductile. The lattice structure for ions is a regular regular pattern but quite quite rigid. A sufficient force will cause the layers to split apart resulting in a brittle nature. iii. 

Electrolysis

Electrolysis experiments show that the ions can be separated using an electrical current. (Refer to fig 1.4.8 p83) iv. 

Electron density maps

Measurements from electron density maps are able to pinpoint the positions of ions, along with their size and charge density. (Refer to fig 1.4.10 p83) 7. 

Lattice energy

The formation of an ionic lattice involves a release in energy so has a negative enthalpy. By compariso comparison, n, bond formation requires energy so these have positive enthalpies. Lattice energy can be calculated using the Born-Haber cycle. (Refer to HSW box p84) a. 

What affects lattice energies?

Lattice energy is affected by both the size and charge of ions. The lattice energy becomes less negative as the size of the ions increases. The lattice energy becomes more negative as the charge on the ions increases. The electrostatic force of attraction attrac tion can be calculated using Coulomb’s law.  law.  (Refer to equation p85)

 

b. 

Predicting stability

The more exothermic a compounds lattice energy is, the more stable it is. Theoretical lattice energies can be used to make predictions about the stability of less familiar elements. (Refer to fig 1.4.12 p86) c. 

Polarisation in ionic bonds

Ionic bonds can be distorted by the attraction of the positive cation to the negative outer electrons electrons of the anion. If the distortion is great great enough, th the e bond can begin to exhibit some covalent character. The polarising power of the cation depends on the charge density which is linked to the charge and and the ionic radius. Cations with a sma smallll ionic radius 3+ have a greater polarising power e.g. Al . The larger an anion is, the more likely it is to be polarised. (Refer to fig 1.4.13 p87) When theoretical and experimental lattice energies are compared, some do not have close agreement for for the same compou compound. nd. If there is a small difference in electronegativities there is more chance of electron sharing and therefore an increase in covalent character. The theoretical theoretical model assu assumes mes that the charge is evenly distributed distributed across an ion, and that all ions are spherical and separate. 8. 

Covalent bonding

Covalent bonds bonds are generally formed formed between non-metals. non-metals. The bonds are formed through the sharing of electrons. The covalent bond is a balance between the attractive force pulling the two nuclei together (due to the electron density between the nuclei) and the repulsion between two two positively charg charged ed nuclei. This balanced dist distance ance is the bond length. The amount of energy required require d to form the bond is the bond enthalpy. 9. 

Dot and cross diagrams

These diagrams show how the outer electrons are shared between atoms, but they do not show the actual positions of bonding atoms atoms in the molecule. (Refer to fig 1.4.16 p89)

 

10. 

Dative covalent bonds

Sometimes both electrons in a covalent bond come from the same atom. This is known as a dative covalent bond or a coordinate bond. Example 1 The ammonium ammonium ion

Example 2 Aluminium chloride dimers

These bonds are often found in oxides, e.g. carbon monoxide. 11. 

Evidence for the nature of covalent bonds

(Refer to HSW box p90) a. 

Molecule sizes

These are strong bonds, with most covalent compounds existing as simple molecules. These have have low melting and boiling boiling points. points. However, there are some exceptions e.g. carbon allotropes, silicon (IV) oxide SiO2. These have high high melting melting and boiling points. points. b. 

Electron density maps and the shapes of molecules

Electron density maps show that covalent bonds are directional towards areas of electronegativity. electronegativity. Lone pairs of ele electrons ctrons (pairs of unshared unshared electrons) also form dense areas of electronegativity. electronegativity. Each area of electronegativity electronegativity repels the others, although lone pairs tend tend to have a stronger repulsion. This gives covalent molecules a very definite shape within a 3D spatial arrangement. (Refer to fig 1.4.20 p91)

 

Some of the more common molecule shapes are; Shape of molecule Linear

Example

Bond angle

No of lone pairs

Trigonal planar

Tetrahedral

 V-shaped

NH3

12. 

Metallic bonding

Metals are good conductors conductors of heat and electricity. There is a strong attraction between the positive ions and the ‘sea of delocalised’ electrons. (Refer to fig 1.4.21 p93) a. 

Properties of metals

(Refer to HSW box p92) i. 

Conducting electricity

 All metals conduct conduct electricity well. The electrons are able able to flow through the the structure to ‘carry’ the current.  current.  ii. 

High thermal conductivity

The electrons can transmit kinetic energy rapidly through the lattice structure. The movement of electrons with a high kinetic energy e nergy is random so the energy can be transferred to cooler regions of the metal.

 

iii. 

High melting & boiling temperatures

Forces between the metal ions are large, so require a lot of energy to overcome them. The delocalised electrons also act as as a ‘glue’ to hold the structure together. iv. 

Malleable & ductile

They can be hammered hammered into shape and drawn out out into wires. The lattice structure, unlike in ionic ionic compounds, compounds, has some flexibility. The ions are able to move their positions slightly with the movement of the delocalised electrons. b. 

Using metals

Copper was used for the the first time abou aboutt 10 000 years ago. In the Middle East steel was produced about 4000 years ago, but it was about 2500 years ago when this began in Europe. The properties of metals make them useful and practical for a variety of different uses. (Refer to fig 1.4.22 p93)

 

Unit 1.5 Introductory Introducto ry Organic Chemistry (p94-107 in text book) 1.  What is organic chemistry? Organic chemistry is considered considered to be the study of carbon chemistry. chemistry. Carbon forms a vast number of compounds due to its ability to form single, double and triple bonds. (Refer to HSW box p94) 1807 – Jons 1807 –  Jons Jacob Berzelius observed that chemicals could be divided into two groups depending depending on their behaviour. behaviour. He classified organic organic molecules as those that would burn burn or char on heating heating.. At this time, scientists scientists believed that organic molecules could only come from living things and that they couldn’t be synthesised (produced artificially).  artificially).  1828 – Friedrich 1828 –  Friedrich Wohler demonstrated demonstrated synthesised synthesised ammonium ammonium cyanate. He demonstrated that this was the same as the urea found in urine. 1846 – Christian 1846 –  Christian Schonbien accidentally combined nitric and sulphuric acids with cellulose. This produced produced nitrocellulose, nitrocellulose, an unstable unstable explosive. 2. 

The vast range of organic compounds

Carbon forms around around 7 million differen differentt compounds. Some of these are complex molecules that make up living cells, enzymes, plastics etc. (Refer to HSW box p95) Many drugs are organic compounds, and do actually come from plant or animal extracts. E.g. aspirin

 

3.  The difference between hazard and risk Many organic substances require special handling. a. 

Hazard

This is essentially the problem associated with handling the chemical e.g. toxic, corrosive etc b. 

Risk

This is the chance that it will actually cause harm. It is dependant upon precautions such as the volumes being used, the expertise of the person handling the chemical, the safety equipment available. available. c. 

Ways of reducing risk

(Refer to p97 & 98) i.  ii.  iii.  iv. 

v. 

Working on a smaller scale – scale – it  it is easier to contain the reaction e.g. in a fume cupboard Taking specific precautions or using alternative techniques – techniques – the  the lowest concentration for a given chemical that works in the reaction can be used rather than more harmful concentrations Careful use of safety measures – measures  – using  using equipment such as goggles and fume cupboards Changing the conditions under which a reaction takes place – place –   this could be to lower the temperature temperature to reduce excessive fumes. The equilibrium position may be changed but the same products are given off Using less hazardous substan substances ces –  – in  in some reactions alternative substances can be used to reduce the hazards involved. CLEAPPS provides a list of suitable alternatives for many reactions

Pesticides are chemicals whose use must be monitored and controlled due to the potential hazards hazards associated associated with their their use. DDT is one such chemical chemical that resulted in toxic build up in food chains. (Refer to HSW box p98 & 99)

 

4.  The properties of the carbon atom Carbon forms 4 covalent bonds. Carbon atoms are unique in their ability to form covalent bonds with carbon and other non-metal atoms at the same time. It can also form single, double or triple bonds with itself. Consequently, it can form a large array of different molecules. 5.  Classifying organic compounds There are millions of different organic molecules that can be classified in a number of different ways. a.  i.  ii.  iii.  b. 

Carbon chain arrangement  Aliphatic molecules molecules have straight straight or br branched anched chain carbon skeletons e.g. alkanes  Alicyclic molecules molecules consist of closed ring structures e.g. cyclobutane cyclobutane  Arenes are all derived derived from th the e benzene molecule molecule (covered in A2) Functionall groups Functiona

There are a variety of functional groups that can be formed within organic molecules. Each homologous homologous series has a general general formula. formula. The molecules molecules within the series will have similar chemical properties due to the functional group but will have varying physical properties e.g. boiling points. (Refer to table 1.5.2 p102)

 

  Homologous series

General formula

Functional group

Example

 

6.  Representing organic compounds There are several different ways of representing organic molecules, depending on the level of detail required. (Refer to p103) E.g. propan-1-ol Empirical formula

Molecular formula

Structural formula

Displayed formula

Skeletal formula

7.  Nomenclature (naming molecules) The IUPAC (International Union of Pure and Applied Chemistry) system is used in naming organic molecules. a. 

Prefix

The prefix of an aliphatic compound relates to the number of carbon atoms that make up the longest chain within the molecule. (Refer to table 1.5.3 p104)

 

Prefix Meth-

b. 

Number of carbon atoms in the main carbon chain 1

Suffix

The suffix of the name relates directly to the main functional group found in the molecule. (Refer to table 1.5.4 p104) Suffix -ane

Number of carbon atoms in the main carbon chain alkane

 

  c. 

Branching

Organic molecules often have branches. The alkyl groups that that form the side chains have the general formula C nH2n+1  E.g. (Refer to p105)  Alkyl group methyl

Formula CH3-

ethyl propyl butyl pentyl hexyl

The position of the side chain c hain is indicated by the lowest possible number e.g. 3-ethylhexane

If more than one side chain is present, these are named in alphabetical alphabetical order e.g. 3-ethyl-2-methylpentane 3-ethyl-2-methylpentane

 

If the side chains are on the same carbon atom the numerical position is repeated e.g. e .g. 2,2-dimethylpentane 2,2-dimethylpentane

d. 

Isomerism

Isomers share the same molecular formula but have a different arrangement of atoms. There are several different different types of isome isomerism, rism, some of which won’t be covered in this unit.  unit.   i. 

Structural isomerism

There are 3 types of structural isomers. Type of structural

Examples

isomer Chain

Positional

Functional group

Butane

2-methylpropane 2-methylpropane

Propan-1-ol

Propan-2-ol

Propanal

Propanone

There are some pitfalls in drawing isomers which must be avoided; kink in the chain rotated molecule (Refer to p107)

 

ii. 

Stereoisomerism

These are molecules that are non-super imposable mirror images of each other. This arises when the the 3D arrangement of bonds bonds in the molecule allow allow different possible orientations of the atoms.

 

Unit 1.6 The Alkanes  A Family of Saturated Hydrocarbons Hydrocarbons  – 

(p108 –  123  123 in text book) 1. 

Hydrocarbons

Hydrocarbons Hydrocarbo ns are molecules made from hydrogen and carbon only. Several different homologous homologous series ma make ke up the the hydrocarbons. They are all insoluble in water and given sufficient oxygen, will combust to produce carbon dioxide and and water. The 3 families families are alkanes, alkenes and alkynes. 2. 

General properties of the alkanes

These are saturated hydrocarbons, hydrocarbons, so have the maximum amount of hydrogens bonded bonded to the carbon atom atoms. s. The general formula formula is CnH2n+2 and these molecules exist as both straight chained and branched structures. (Refer to fig 1.6.1 p108) The boiling points of alkanes increase as the molecular formula increases i.e. as more carbon atoms are added. However, straight chain molecules have higher boiling points than an isomer that is branched. The straight chain molecules molecules are able to exist clos closer er to each other and as a result re sult have stronger intermolecula intermolecularr forces of attraction to overcome. 3. 

Where are alkanes obtained from?

(Refer to fig 1.6.2 p109) The most important important source of alkanes is the fossil fue fuels. ls. Oil was formed millions of years ago in the seas. Tiny plants and animals fell to the bottom and over time layers formed formed above them. them. Compression in the the absence of oxygen caused the formation formation of crude oil and natural natural gas. The oil and gas become trapped under non-porous rock which must be drilled through to access the fuel. Coal was formed in a similar way from plant remains (mainly large ferns) on land. 4. 

The economic importance of crude oil

(Refer to HSW box p110) Modern life depends on crude oil which provides fuel for transport and generating electricity. electricity. It also provides provides many of the raw materials materials needed in

 

the chemical industry. industry. However, it is a finite resource and most most deposits are are found in politically sensitive areas of the world. Until the 1970s oil was regarded as a cheap resource. resource. Then the Middle Middle Eastern countries supplying the oil decided to produce less and increase their prices. Since then, oil prices have have varied greatly ((even even on a daily basis). However, the general trend trend has been an increase. increase. This has a knock on effect effect on users of oil, e.g. transport costs have increased. 5. 

Using crude oil

The crude oil extracted from the the ground is of little us use e without processin processing. g. It is a mixture of many different different hydrocarbons. It undergoes primary primary distillation distillation (fractional distillation on an industrial scale) to separate the different fractions. It can then un undergo dergo further fractional distillation distillation to give give a pure yield of a particular alkane. (Refer to fig 1.6.4 p111) Fraction

Percentage

Length of carbon chain

Use(s)

Refinery gas

Gasoline

Kerosene

Diesel oil / Gas oil

Residue

6. 

Different fractions in demand

Different reserves of crude oil have the different fractions in varying quantities. The lighter fractions fractions are in much higher higher demand than the heavier heavier fractions, which can make up to 50% of the crude oil mixture. (Refer to fig 1.6.5 p112)

 

Cracking is used to break down the longer carbon chains into the alkanes found in the lighter fractions along with some short chain alkenes such as ethane, which are also extremely useful. Cracking requires extremely high temperatures, so catalysts are often used. This process is thereby known known as catalytic catalytic cracking. The catalysts used are are often crystalline aluminosilicates (zeolites). (Refer to fig 1.6.6 p112) 7. 

Knocking and the need for catalytic reforming

(Refer to HSW box p113) The alkanes in petrol are generally generally 5-10 carbon atoms atoms in length. The power is produced by the explosive combustion in cylinders. The smooth running of the car depends on the explosion happenin happening g at exactly the right time. If the explosion ha happens ppens too early ‘kno ‘knocking’ cking’ occurs which produces a loss of power. power. This is usually the the case when the petrol contains contains a large proportion of straight straight chain alkanes alkanes since they easily ignite. ignite. Petrol with a high proportion of branched chain alkanes are more efficient fuels since they take longer to ignite. Fuels are given an octane rating to indicate the ratio of straight chain to branched chain molecules. molecules. A fully straight straight chain fuel has a rating rating of 0, while a fully branched chain fuel has a rating of 100. Fuel pumps show the octane rating by the Research Octane Number (RON). Ordinary unleaded is 95 while super unleaded is 98. Most modern cars have high-performance engines so require fuels with a high octane rating. Tetraet Tetraethyl hyl lead(IV) used to be used to prevent knocking but resulted in lead pollution. pollution. Nowadays, catalytic reforming is used to produce produce branched chain molecules from straight chain molecules in a similar process to catalytic cracking. 8. 

The chemical properties of the alkanes

 Apart from combustion, combustion, the the alkanes ar are e very unreactive. This makes them them very useful. They are non-corrosive non-corrosive with m metals etals making them good lubricating oils. They are also harm harmless less to the skin e.g. petroleu petroleum m jelly.  At room temperature temperature the alkanes alkanes are un unaffected affected by concentrated concentrated acids and and alkalis. They are not affected by oxidising agents or reactive metals. The bonds in alkanes involve a very even sharing of electrons since the electronegativities electronegativiti es of carbon and hydrogen are very close.

 

 Almost all the reactions with alkanes alkanes occur due due to the forma formation tion of free radicals, which contain contain an unpaired electron. They have high activation activation energies, but once overcome the reaction proceeds quickly in the gas phase. 9. 

Breaking bonds

This is known as bond fission. Since a covalent bond involves involves the sharin sharing g of 2 electrons, there are 2 ways in which the electrons can be shared out when a bond is broken; a. 

Homolytic fission

This involves an equal equal sharing of the once bon bonding ding electrons. As a result, each atom in the bond receives an unpaired electron to form a free radical. These free radicals are are extremely reactive. reactive. The reaction of one free radical radical with a substance usually results in the formation of another free radical. (Refer to fig 1.6.8 p114) Free radicals in the body are now thought to be carcinogens. carcinogens . The number of free radicals in the body can be reduced by cutting down on the intake of food high in free radicals – radicals – burnt  burnt toast and charred food from the BBQ for example. Fresh fruit and vegetables vegetables rich in vita vitamins mins A, C and E help the the enzyme superoxide dismutase to deactivate free radicals to prevent any damage they may cause. (Refer to HSW box p115) Free radicals are also linked linked to the aging process. process. The theory is that that they attack the cross links between collagen fibres in the skin causing it to appear less flexible. They also attack the DNA in nuclei in body cells which is thought to be the cause of cancer.  Antioxidants are are viewed to ‘mop ‘mop up’ the the free radicals in the body. Vitamin E is a well known example of a natural free radical inhibitor. (Refer to HSW box p119) b. 

Heterolytic fission

This involves an unequal unequal sharing of the once bonding bonding electrons. This results in the formation of ions. (Refer to fig 1.6.9 p115)

 

10. 

The reactions of alkanes

There are just 2 common types of reactions that alkanes undergo a. 

Cracking

When heated to high temperatures in the absence of air alkanes will undergo cracking. This is a thermal decomposition decomposition react reaction ion to form smaller smaller molecules. b. 

Combustion

When heated in a plentiful plentiful supply of air, comb combustion ustion will take place. place. This is a highly exothermic reaction. reaction. Gaseous alka alkanes nes will burn completely but but solid and liquid fuels are difficult to ignite. This reaction is highly important since it is used to generate electricity, fuel fires in the home, provide central heating, cooking and transport. (Refer to fig 1.6.12 p116) Methane, propane propane and bu butane tane are all commonly u used sed as fuels. Where possible, natural gas is piped into homes. When this can’t happen, canisters canister s of propane can be used. Propane can be readily liquefied making it easy to store and transport. transport. It won’t burn until until it returns to the ga gaseous seous state making it safe to use. Incomplete combustion can produce carbon monoxide and has resulted re sulted in fatalities. Regular maintenance maintenance of gas appliances appliances is required to reduce reduce the risk of CO poisoning. poisoning. There are strict regulations regulations imposed on landlords to ensure the safety of tenants. (Refer to HSW box p117) c. 

Reaction with chlorine

This reaction is known as free radical substitution, substitution, where chlorine radicals substitute for hydrogen atoms on an alkane. The reaction requires requires a large input input of energy in the form of UV light. light. This provides the energy to break the bond in chlorine molecules and is known as initiation. It is an example of homolytic homolytic fission.

In propagation propagation the chlorine radical reacts with the the alkane. At the end, another free radical is generated. generated. This process can be be repeated hundreds hundreds of times as an explosive chain reaction.

 

 

(Refer to fig 1.6.14 p118) In termination, termination, 2 radicals combine combine in a highly highly exothermic exothermic reaction. This happens every few thousand reactions.

(Refer to fig 1.6.15 p118) Provided there is an excess of chlorine, further substitutions substitutions can take place. Two of these compounds compounds are well known. Trichloromethane Trichloromethane (chloroform) CHCl3 was one of the first anaesthetics anaesthetics used. used. Tetrachlorometh Tetrachloromethane ane CCl4 was widely used as a solvent until its carcinogenic properti properties es were recognised. 11. 

Ethical issues

(Refer to HSW boxes p120 –  123)  123) a. 

Cars and society

One of the major uses of alkanes alkanes is as fuel for transport. transport. Cars are one of the the major causes of pollution pollution on the planet. planet. Carbon monoxide is formed formed in the engine which is a greenhouse greenhouse gas as well as being h highly ighly toxic. Oxides of nitrogen and sulphur are also produced which cause acid rain and smog. Catalytic converters convert CO and NO x into less harmful gases but cannot remove the CO2. b. 

The greenhouse effect

Greenhouse gases include water vapour, CO 2 and methane, 2 of which are products of hydrocarbon hydrocarbon combu combustion. stion. These gases cause the temperature temperature on Earth to rise due to re-radiation of energy from the Sun. c. 

Developing new fuels

 

Biofuels such as ethanol, along with hydrogen powered and solar powered vehicles vehicl es have been looked at as alternatives. alternative s. However, the amount of land needed to grow the plants for biofuels is vast. d. 

Battery-powered vehicles

Electricity is an alternative, but this is often produced through the combustion of fossil fuels. The technology is st still ill in its early stages stages and the size of the battery required is large and heavy providing far less energy than petrol. Recharging can then take up to 12 hours. e. 

The hydrogen cell

Electrolysis of hydrogen hydrogen takes place in a cell to form wa water ter and heat. heat. The hydrogen cell motorbike can travel at around 50mph for 100 miles at a cost of £2. However, this technology technology is in its early stages. Petrol stations stations aren’t aren’t yet equipped to offer refuelling of hydrogen, even though it is relatively safe.

 

Unit 1.7 The Alkenes  A Family of Unsaturated Hydrocarbons  – 

(p124 –  141  141 in text book) 1. 

The carbon  carbon double bond  – 

These are unsaturated unsaturated hydrocarbons due to the existence of the double bond, which makes them them more reactiv reactive. e. There general general formula is CnH2n  In a single bond (sigma ơ bond), bond), the electron cloud is symmetrical about the central axis of the molecule. In a double bond not only does the ơ bond exist, but so does a pi π bond. The electron density is concentrated in 2 regions on either side of the bond axis. Consequently, there there is no rotation rotation around around the double double bond. (Refer to fig 1.7.2 p124) The π bond is reactive towards electrophiles, providing an electron pair to bond with electron seeking seeking species. It also results in geometric geometric isomerism due to the lack of rotation around the double bond. 2. 

Ethene

(Refer to HSW box p125) Ethene is a plant hormone hormone used in ripening ripening fruit. It is used on a commercial commercial basis to ripen fruit ready ready for supermarket supermarket shelves. However, this could could be having adverse affects on the planet due to the many miles travelled in transporting some fruit across the globe. 3. 

Geometric isomerism

These have different physical properties e.g. boiling point, but identical chemical properties. properties. They are a form of stereoisomers stereoisomers since there are are 2 possible arrangements of the species on the double bond. The traditional way of naming these is cis-trans cis -trans isomerism e.g. Cis-but-2-ene

Tran-but-2-ene

 

However, this method method of naming has its limitations. The IUPAC system is now the E-Z isomerism isomerism system system.. This is capable of working for a allll geometric isomers. Groups around the double bond are ranked on their atomic number. The atom with the highest number has the highest priority. If the two groups with higher priority are on the same side they are zusammen , the Z-isomer. If the higher priority groups are on opposite sides they are entgagen , the Eisomer. E.g. 1-bromo-2-chloro-2-fluoro-1-iodoethene Z-isomer

E-isomer

The body is sensitive sensitive to geometric isomers. isomers. This has been studied in cooked tomatoes. (Refer to HSW box p127) 4. 

Reactions of the alkenes

 Alkenes are more more reactive than than alkanes due due to the presence of the the double bond. The high electron density density of the double bo bond nd means they are at attacked tacked by electrophiles. An electrophile electrophile is an electron deficient species that that can form + a covalent bond e.g. H  ion a. 

Electrophilic addition reactions

i. 

Reaction with hydrogen

 Alkenes can be reduced to alk alkanes anes through through addition addition of hydrogen at around 200°C with Ni as a catalyst.

 

ii. 

Reaction with the halogens

 A halogen molecule molecule can be added added across th the e double bond bond to produce produce a dihalogenoalkane dihalogenoalk ane e.g. ethene reacting with with chlorine. This is possible possible due to the temporary dipole in the halogen molecules.

iii. 

Testing for alkenes with bromine water

The major product is 2-bromoethanol but some 1,2-dibromoethane is also formed. Bromine water water is decolourised decolourised in a positive positive test.

iv. 

Reaction with hydrogen halides

This reaction proceeds rapidly at room temperature due to the permanent dipole in the hydrogen halide halide . The resulting compound compound is a halogenoalkane. halogenoalkane. In asymmetric alkenes, there are 2 possible products, but not in equal proportions. The major product product can be predicted using Markovnikov’s Markovnikov’s rule:  rule:  When HX adds across an asymmetric double bond, the major  product formed is is the molecule in which hydrogen hydrogen adds to the the carbon atom in the double bond with the greater number of hydrogen atoms already attached to it.

 

v. 

Reaction with acidified potassium manganate(VII) manganate(VII)

This is an oxidation reaction. reaction. The potassium potassium manganate(VII) manganate(VII) is decolourised in the production of a diol. This reaction can also be used to distinguish between alkenes and alkanes.

It is a useful product, but is produced industrially industrially from epoxyethane. vi. 

Reaction mechanism

This is an electrophilic addition addition mechanism. mechanism. A carbocation (carbonium (carbonium ion) is formed as an intermediate. intermediate. Methyl groups groups can donate donate electron density to stabilise the positive ion (this is known as an inductive effect). E.g. Ethene reacting with hydrogen bromide

 

b. 

Polymerisation

 A polymer is a very large m molecule olecule made up of long cha chains ins of smaller smaller units (monomers) joined together. together. Artificial polym polymers ers include plastics and and are important materials for modern life.  Alkenes undergo addition poly polymerisation merisation where the double double bond breaks breaks open to link together repeating units in a long chain. E.g. polythene

Polypropene

Polychloroethene Polychloroethen e (PVC)

Polytetrafluoroethene

 

  5. 

The properties of polymers

The properties of polymers depend on the chains of monomers from which they are formed. They can vary greatly. greatly. Feature

Property

The average length of the polymer chain

Branching of the chain

The presence of intermolecular forces between chains Cross-links between chains

6. 

Polymer problems and solutions

(Refer to HSW box p136 –  137)  137) a. 

Problems

The production of synthetic polymers has saved many natural materials from destruction, but there is a down side to this. i. 

Energy costs

Polymers are produced produced from fossil fuels. There are hidden energy co costs sts to this. Around 4% of the worlds worlds fossil fuel supply is used used to generate the electricity needed for polymer production. ii. 

Resources used

There is a limited supply of fo fossil ssil fuels with rising costs. This gives an incentive to produce polymers from different monomer units.

 

iii. 

Disposal problems

They aren’t easy to dispose dispose of and cause a substantial substantial waste problem. problem. A variety of toxic gases are released when burnt, including hydrogen cyanide. They also cause danger to wild animals. iv. 

Carbon footprint

The use of fossil fuels in polymerisation releases carbon thereby increas increasing ing our carbon footprint. footprint. The quantity of non-biodegradable non-biodegradable polymers is greatly increasing. b. 

Solutions

i. 

Renewable energy sources

These could be used to generate the electricity required to make polymers. ii. 

Reducing use of polymer products

There is a move to reduce the the amount of packa packaging ging used e.g. in food. Higher quality plastics are being used in packaging to encourage reuse e.g. ‘bag e.g. ‘bag for  for  life’.   life’. iii. 

Recycling

Many plastics are difficult to dispose of in a way that doesn’t damage the environment. Many sy synthetic nthetic polymers polymers are non-biodegradable. non-biodegradable. iv. 

Recycling thermoplastics

So far only thermoplastics (plastics that soften on heating) can be recycled. recyc led. If the plastic that is recovered after recycling is as good as the original, then effectively all the hidden energy from its original production has been saved apart from the energy used in recycling.

 

7. 

How are plastics recycled?

There are 2 main ways of recycling; Type of recycling Mechanical recycling

Key points

Chemical recycling

8. 

Biodegradable polymers

Biopolymers are now being produced through the modification of natural polymers such as cellulose an and d starch. Bacterial fermentation fermentation is used tto o break down the polymer. The degradation may take 20-30 years but the product will eventually be destroyed. (Refer to HSW box p139) 9. 

Energy recovery

Burning waste polymer products can be used to generate electricity, reducing the need for fossil fossil fuels. The incinerators used used have special special pollution pollution controls. Pyrolysis and gasification are both processes where polymer wastes are converted into energy-rich fuels. Many of these processes aren’t very efficient at the moment. 10. 

Life cycle analysis

(Refer to fig 1.2.25 p140) This is used to to quantify the effect of polymers polymers on the environmen environment. t. An inventory is made of the materials and energy used as well as environmental emissions. (Refer to table 1.7.5 p141)