Atoms, Elements and Compounds

ATOMS, ELEMENTS AND COMPOUNDS Some atoms are very reluctant to combine with other atoms and exist in the air around us as single atoms. These are the Noble Gases and have very stable electron arrangements e.g. 2, 2,8 and 2,8,8 because their outer shells are full. The first three are shown in the diagrams below and explains why Noble Gases are so reluctant to form compounds with other elements.

(atomic number) electron arrangement All other atoms therefore, bond together to become electronically more stable, that is to become like Noble Gases in electron arrangement. Bonding produces new substances and usually involves only the 'outer shell' or 'valency' electrons and atoms can bond in two ways. The phrase CHEMICAL BOND refers to the strong electrical force of attraction between the atoms or ions in the structure. The combining power of an atom is sometimes referred to as its valency and its value is linked to the number of outer electrons of the original uncombined atom (see examples later). Each type of chemical bonding is VERY briefly described below, with links to more detailed notes.

IONIC BONDING Ionic bonds are formed by one atom transferring electrons to another atom to form ions. Elements consist of neutral atoms or molecules, the electrical neutrality is because the number of positive protons equals the number of surrounding negative electrons in their respective energy levels. Charged particles called IONS are atoms, or groups of atoms, which have lost or gained one or more electrons to have a overall net electrical positive charge or negative charge. In losing or gaining electrons, the atoms try to attain a stable electron arrangement of a noble gas e.g. a full outer shell of electrons. For a given atom, a nearly full shell will try to gain electrons and a nearly empty shell will tend to lose electrons The atom losing electrons forms a positive ion (a cation) and is usually a metal. The overall charge on the ion is positive due to excess positive nuclear charge (protons do NOT change in chemical reactions) e.g. Group 1 alkali metals lose their single outer electron to form single positive ions e.g. Na ==> Na+ + 2e– Group 2 metals lose their two outer electrons to form doubly charged positive ions e.g. Mg ==> Mg2+ + 2e–

The atom gaining electrons forms a negative ion (an anion) and is usually a non–metallic element. The overall charge on the ion is negative because of the gain, and therefore excess, of negative electrons e.g. Group 7 halogen atoms gain one electron to form a singly charged negative ion e.g. Cl + e– ==> Cl– Group 6 non–metals gain two electrons to form a doubly charged negative ion e.g. O + 2e– ==> O2– Therefore an IONIC BOND IS THE FORCE OF ATTRACTION BETWEEN ADJACENT IONS OF OPPOSITE CHARGE. Which electronic structures are the most stable? because this what atoms will try to get to electronically!

symbol (atomic number) electron arrangement When atoms LOSE OR GAIN ELECTRONS, they try to attain the electron structure (electron configuration) of the electronically very stable atoms of the Group 0 Noble Gases eg helium (2), neon (2.8) or argon (2.8.8) etc. quite simply because these are the most stable electron arrangements with a full outer shell of electrons (full highest energy level). In advanced level chemistry you will encounter examples of electronic structures of ions that are NOT those of a Noble Gas. Its a good idea to have some idea of where the elements are in the periodic table, and their electronic structure, before looking at the theoretical electronic model for ionic bonding in ionic compounds

The black zig–zag line 'roughly' divides the metals on the left from the non– metals on the right of the elements of the Periodic Table.

The electronic structures of the first 20 elements of the Periodic Table You need to know about these to understand the details of ionic chemical bonding P m Pm d ea e trt t ao a lf l st s h e= m> o dn eo rn n–

Pm ee ri t oa di l cs T a bl e P d = p e Gri G G G po ppp 12d,345670 G p = g r o u p

1

H

Note that H does

2

1not H readil e y fit

into any grou p

ato 1 34 mi 567890

2L B B C N O F N iec nu e 311mb111111 er 12 345678

CA N M S P S C A agh lilr e m ic al S y m b

ol e g 2222222223333333 12 0124567890123456 93

4 C S TC M F C N C Z G A S B K KV acirn eo iu n aeser 3344444444445555 3 5 7801234567890124 9 3

5 R SZ N M T R P A C ISTX Y I b rrb o cu h d g d nb ee T r a n 55si 888888 ti 656 123456 C Bo T P B P A R san lb io tn M et al s Gp 1 Alk ali Me tal s Gp 2 Alk ali ne Ear th Me tal s Gp 7 Hal og en s Gp 0 No ble Ga ses Ch em

ical bo ndi ng co m me nts ab out the sel ect ed ele me nts hig hli ght ed in whi te e.g . Wh en the me tal s on the left co mb ine wit h the no n– me tal s on the rig ht, an ion ic bo nd is for

me d e.g . the for ma tio n of an ion ic co mp ou nd like so diu m chl ori de Na Cl

All the atoms of the metallic elements on the left–hand side of the Periodic Table eg (Groups 1/2) have 1/2 electrons in their outer shell (highest energy level) which are readily lost to form a positive ion of charge +1/+2 (cations) eg sodium, potassium, magnesium and calcium etc. The electronic structure of these stable positive ions are those of a Noble Gas with a full outer shell. eg Na [2.8.1] ==> Na+ [2.8] like neon + e–, or Ca [2.8.8.2] ==> Ca2+ [2.8.8] like argon + 2e– The atoms of the non–metallic elements on the right–hand side of the Periodic Table eg (Groups 6/7) have 6/7 electrons in their outer shell and try to gain 2/1 electrons to become electronically stable like a Noble Gas with a full outer shell of electrons eg oxygen and sulfur in Group 6 and the Group 7 Halogens – fluorine, chlorine, bromine and iodine. The electronic structure of these stable negative ions are those of a Noble Gas with a full outer shell. eg O [2.6] + 2e– ==> O2– [2.8] like neon, or Cl [2.8.7] + e– ==> Cl– [2.8.8] like argon

Example 1: A Group 1 Alkali Metal + a Group 7 Halogen non–metal e.g. sodium + chlorine ==> sodium chloride NaCl or ionic formula Na+Cl– In terms of electron arrangement in the formation of the ionic compound sodium chloride, the sodium donates its outer electron to a chlorine atom forming a single positive sodium ion and a single negative chloride ion. The atoms have become stable ions, because electronically via electron transfer ...

... sodium becomes like neon (sodium ion, Na+) and chlorine like argon (chloride ion, Cl–). Na (2.8.1) + Cl (2.8.7) ==> Na+ (2.8) Cl– (2.8.8) can be summarised electronically to give the stable 'noble gas' structures as [2,8,1] + [2,8,7] ==> [2,8]+ [2,8,8]– so both the sodium and chloride ions have a full outer shell like a noble gas

ONE

atom combines with

ONE

atom to form

Note in this electron diagram, only the original outer electrons are shown above. The outer electron of the sodium atom (2.8.1) is transferred to the outer shell of the chlorine atom (2.8.7) giving it a complete octet shell of outer electrons, just like a noble gas (2.8.8). At the same time, the sodium ion also attains a stable noble gas electron structure (2.8). The valencies of Na and Cl are both 1, that is, the numerical charge on the ions. sodium fluoride NaF, potassium bromide KBr and lithium iodide LiI etc. will all be electronically similar. Only the outer valency electrons of the chloride ion are shown, the 'blob' electron represents the electron from the sodium atom which is accepted by the chlorine atom to form the chloride ion. The charge on the sodium ion Na+ is +1 units (by convention shown as just +) because there is one more positive proton than there are negative electrons in the sodium ion (11p, 10e). The charge on the chloride ion Cl– is –1 units (by convention shown as just –) because there is one more negative electron than there are positive protons in the chloride ion (17p, 18e). Note:

would represent the full electronic structure diagram of the sodium ion [2.8] and the chloride ion [2.8.8], hence the full electronic structure of sodium chloride. Note that the 'blob' and 'x' electrons are identical, but their use is just a useful visual device to show how the ions are formed. The blue circle represents the nucleus. The electronic dot & cross Lewis diagram for the ionic bonding in sodium chloride

Li is 2.1, K is 2.8.8.1, F is 2.7, rest of dot and cross diagram is up to you. Gp1\ 7

F

Cl

Br

I

Li

LiF LiCl LiBr LiI

Na

Na F

K

KF KCl KBr

KI

Rb

Rb

Rb Cl

Rb I

Cs

Cs F

Cs Cl

Na Cl

NaB Na r I

Rb Br

CsB Cs r I

All the formula highlighted in yellow can be described in the same way as sodium chloride The Group 1 Alkali Metal atom loses one electron to form a singly charged positive ion The Group 7 Halogen atom gains one electron to form a singly charged negative ion

Example 2: A Group 2 Alkaline Earth Metal + a Group 7 Halogen non–metal e.g. magnesium + chlorine ==> magnesium chloride MgCl2 or ionic formula Mg2+(Cl–)2 In terms of electron arrangement in the formation of the ionic compound magnesium chloride, the magnesium donates its two outer electrons to two chlorine atoms forming a double positive magnesium ion and two single negative chloride ions via electron transfer. The atoms have become stable ions, because electronically, magnesium becomes like neon and chlorine like argon. Mg (2.8.2) + 2Cl (2.8.7) ==> Mg2+ (2.8) 2Cl– (2.8.8) can be summarised electronically as [2,8,2] + 2[2,8,7] ==> [2,8] 2+ [2,8,8]–2 via electron transfer so both the magnesium and chloride ions have a full outer shell of electrons like a noble gas

ONE

atom combines with

TWO

atoms to

form Note in this electron diagram, only the original outer electrons are shown above.

The outer electrons of the magnesium atoms (2.8.2) are transferred to the outer shell of the chlorine atom (2.8.7) giving it a complete octet shell of outer electrons, just like a noble gas (2.8.8). At the same time, the magnesium ion also attains a stable noble gas electron structure (2.8). NOTE You can draw two separate chloride ions, but in these examples square brackets and a number subscript have been used, as in ordinary chemical formula. The valency of Mg is 2 and chlorine 1, i.e. the numerical charges of the ions. Beryllium fluoride BeF2, magnesium bromide MgBr2, calcium chloride CaCl2 or barium iodide BaI2 etc. will all be electronically similar.

represents the full electronic structure diagram of the magnesium ion [2.8] and the chloride ion [2.8.8], hence the full electronic structure of magnesium chloride. Note that the 'blob' and 'x' electrons are identical, but their use is just a useful visual device to show how the ion is formed. The blue circle represents the nucleus. The electronic dot & cross Lewis diagram for the ionic bonding in magnesium chloride

Ca is 2.8.8.2, Cl is 2.8.7, F is 2.7 rest of dot and cross diagrams are up to you, but calcium chloride is shown below. The calcium atoms transfer their two outer electrons to the outer shell of two chlorine atoms

calcium chloride The electronic dot & cross Lewis diagram for the ionic bonding in calcium chloride The two outer electrons of the calcium atoms (2.8.8.2) are transferred to the outer shell of two chlorine atoms (2.8.7) until it has a complete octet shell of outer electrons, just like a noble gas (2.8.8). At the

same time, the calcium ion also attains a stable noble gas electron structure (2.8.8). The blue circle represents the nucleus. G p2 F \7 M g

Cl Br

M M M M g g g g Cl Br F2 I2 2

C a

I

2

C C C C a a a a Cl Br F2 I2 2

2

Sr

S Sr Sr S r Cl Br rI F2 2 2 2

B a

B B B B a a a a Cl Br F2 I2 2

2

All the formula highlighted in yellow can be described in the same way as magnesium chloride or calcium chloride The Group 2 Alkaline Earth Metal atom loses two electrons to form a doubly charged positive ion The Group 7 Halogen atom gains one electron to form a singly charged negative ion

Example 3: A Group 3 metal + a Group 7 non–metal e.g. aluminium + fluorine ==> aluminium fluoride AlF3 or ionic formula Al3+(F–)3 In terms of electron arrangement in the formation of the ionic compound aluminium fluoride, the aluminium donates its three outer electrons to three fluorine atoms forming a triple positive aluminium ion and three single negative fluoride ions. The atoms have become stable ions, because aluminium and fluorine becomes electronically like neon via electron transfer. Valency of Al is 3 and F is 1, i.e. equal to the charges on the ions. Al (2.8.3) + 3F (2.7) ==> Al3+ (2.8) 3F– (2.8) can be summarised electronically as [2,8,3] + 3[2,7] ==> [2,8] 3+ [2,8]–3 so both the aluminium and fluoride ions have a full outer shell like a noble gas

ONE

atom combines with

THREE

atoms to

form Note in this electron diagram, only the original outer electrons are shown above. The outer electrons of the aluminium atom (2.8.3) is transferred to the outer shell of the fluorine atoms (2.7) giving them a complete octet shell of outer electrons, just like a noble gas (2.8). At the same time, the aluminium ion also attains a stable noble gas electron structure (2.8).

full electronic structure diagram of aluminium fluoride, the blue circle represents the nucleus. The electronic dot & cross Lewis diagram for the ionic bonding in aluminium fluoride Solid aluminium chloride/bromide/iodide have similar formula but are covalent when vapourised into Al2X6 dimer molecules, but AlCl3 has an ionic lattice in the solid, not sure on solid AlBr3 and AlI3, but these points are best left for an advanced AS–A2 chemistry discussion, not for GCSE students!

Example 4: A Group 1 Alkali Metal + a Group 6 non–metallic element e.g. sodium/potassium + oxygen ==> sodium/potassium oxide Na2O/K2O or ionic formula (Na+)2O2–/ (K+)2O2– In terms of electron arrangement in the formation of the ionic compound sodium oxide, the two sodium/potassium atoms donate their outer electron to one oxygen atom. This results in two single positive potassium ions to one double negative oxide ion via electron transfer. All the ions have the stable electronic structures 2.8.8 (argon like) or 2.8 (neon like). Valencies, sodium/potassium 1, oxygen/sulfur 2. giving the following formulae: Lithium oxide, Li2O, sodium oxide Na2O, sodium sulphide Na2S and potassium K2S etc.

sodium oxide 2Na (2.8.1) + O (2.6) ==> 2Na+ (2.8.8) O2– (2.8) can be summarised electronically as 2[2,8,1] + [2,6] ==> [2,8] +2 [2,8]2– so both the sodium and oxide ions have a full outer shell like a noble gas

TWO

atoms combine with

ONE

atom to form

or

+

+

==>

Note in this electron diagram, only the original outer electrons are shown above. The outer electrons of the sodium atoms (2.8.1) are transferred to the outer shell of the oxygen atom (2.6) until it has a complete octet shell of outer electrons, just like a noble gas (2.8). At the same time, the sodium ion also attains a stable noble gas electron structure (2.8).

full electronic structure diagram of sodium oxide, the blue circle represents the nucleus. The electronic dot & cross Lewis diagram for the ionic bonding in sodium oxide

potassium oxide 2K (2.8.8.1) + O (2.6) ==> 2K+ (2.8.8) O2– (2.8) can be summarised electronically as 2[2,8,8,1] + [2,6] ==> [2,8,8] +2 [2,8]2–

so both the potassium and oxide ions have a full outer shell like a noble gas

TWO

atoms combine with

ONE

atom to form

Note in this electron diagram, only the original outer electrons are shown above. The outer electrons of the potassium atoms (2.8.8.1) are transferred to the outer shell of the oxygen atom (2.6) until it has a complete octet shell of outer electrons, just like a noble gas (2.8). At the same time, the potassium ion also attains a stable noble gas electron structure (2.8.8).

full electronic structure diagram of potassium oxide, the blue circle represents the nucleus. The electronic dot & cross Lewis diagram for the ionic bonding in potassium oxide

The electronic similarities between the two examples are very obvious. Li is 2.1, Na is 2.8.1, S is 2.8.6 (for group 1 sulphide compound), rest of dot and cross diagrams are up to you. e.g. electronic structure diagrams for sodium sulfide Na2S and potassium sulfide K2S

sodium sulfide The electronic dot & cross Lewis diagram for the ionic bonding in sodium sulphide

The outer electrons of the sodium atoms (2.8.1) are transferred to the outer shell of the sulfur atom (2.8.6) until it has a complete octet shell of outer electrons, just like a noble gas (2.8.8). At the same time, the sodium ion also attains a stable noble gas electron structure (2.8).

potassium sulfide The outer electrons of the potassium atoms (2.8.8.1) are transferred to the outer shell of the sulfur atom (2.8.6) until it has a complete octet shell of outer electrons, just like a noble gas (2.8.8). At the same time, the potassium ion also attains a stable noble gas electron structure (2.8.8). The electronic dot & cross Lewis diagram for the ionic bonding in potassium sulphide Gp1\ 6

O

S

Li

Li2O Li2S

Na

Na2 O

Na2 S

K

K2O

K2S

Rb

Rb2 O

Rb2 S

Cs

Cs2 O

Cs2 S

All the formula highlighted in yellow can be described in the same way as sodium oxide, potassium oxide, sodium sulfide or calcium sulphide The Group 1 Alkali Metal atom loses one electron to form a singly charged positive ion The Group 6 non–metal atom gains two electrons to form a doubly charged negative ion

Example 5: A Group 2 Alkaline Earth Metal + a Group 6 non–metallic element e.g. magnesium/calcium + oxygen ==> magnesium/calcium oxide MgO/CaO or ionic formula Mg2+O2–/Ca2+O2– In terms of electron arrangement in the formation of the ionic compound magnesium oxide, one magnesium/calcium atom donates its two outer electrons to one oxygen atom.

This results in a double positive calcium ion to one double negative oxide ion via electron transfer. All the ions have the stable electronic structures 2.8.8 (argon like) or 2.8 (neon like). the valency of both calcium and oxygen is 2. magnesium oxide

ONE

atom combines with

ONE

atom to

form Note in this electron diagram, only the original outer electrons are shown above. The two outer electrons of the magnesium atoms (2.8.2) are transferred to the outer shell of the oxygen atom (2.6) until it has a complete octet shell of outer electrons, just like a noble gas (2.8). At the same time, the magnesium ion also attains a stable noble gas electron structure (2.8).

full electronic structure of magnesium oxide For magnesium oxide: Mg (2.8.2) + O (2.6) ==> Mg2+ (2.8) O2– (2.8) the stable 'noble gas' structures can be summarised electronically as [2,8] + [2,6] ==> [2,8] 2+ [2,8]2– so both the magnesium and oxide ions have a full outer shell like a noble gas The electronic dot & cross Lewis diagram for the ionic bonding in magnesium oxide

calcium oxide Ca (2.8.8.2) + O (2.6) ==> Ca2+ (2.8.8) O2– (2.8) can be summarised electronically as [2,8,8,2] + [2,6] ==> [2,8,8] 2+ [2,8]2–

ONE

atom combines with

ONE

atom to

form Note in this electron diagram, only the original outer electrons are shown above. The two outer electrons of the calcium atoms (2.8.8.2) are transferred to the outer shell of the oxygen atom (2.6) until it has a complete octet shell of outer electrons, just like a noble gas (2.8). At the same time, the calcium ion also attains a stable noble gas electron structure (2.8.8).

full electronic structure of calcium oxide The electronic dot & cross Lewis diagram for the ionic bonding in calcium oxide

Magnesium oxide MgO, magnesium sulphide MgS and calcium sulphide CaS will be similar electronically and give identical giant ionic lattice structures. Group 2 metals lose the two outer electrons to give the stable 2+ positive ion (cation) and S and O, both non–metals in Group 6, have 6 outer electrons and gain 2 electrons to form 2– negative ion (anion). For magnesium sulphide: Mg (2.8.2) + S (2.8.6) ==> Mg2+ (2.8) S2– (2.8.8) For calcium sulphide: Ca (2.8.8.2) + S (2.8.6) ==> Ca2+ (2.8.8) S2– (2.8.8) so both the magnesium/calcium and sulfide ions have a full outer shell like a noble gas The dot and cross (ox) diagrams will be identical to that for calcium oxide above, except Mg instead of Ca (same group) and S instead of O (same group of Periodic Table). eg

electronic structure of magnesium sulfide MgS

electronic structure of calcium sulfide CaS The electronic dot & cross Lewis diagrams for the ionic bonding in magnesium sulphide and calcium sulphide Gp2\ 6

O

S

Mg

Mg O

Mg S

Ca

CaO CaS

Sr

SrO

Ba

BaO BaS

SrS

All the formula highlighted in yellow can be described in the same way as magnesium oxide, magnesium sulphide, calcium oxide or calcium sulphide The Group 2 Alkaline Earth Metal atom loses two electrons to form a doubly charged positive ion The Group 6 non–metal atom gains two electrons to form a doubly charged negative ion

Example 6: A Group 3 metal + a Group 6 non–metal e.g. aluminium + oxygen ==> aluminium oxide Al 2O3 or ionic formula (Al3+)2(O2–)3 In terms of electron arrangement in the formation of the ionic compound aluminium oxide, two aluminium atoms donate their three outer electrons to three oxygen atoms.

This results in two triple positive aluminium ions to three double negative oxide ions via electron transfer. All the ions have the stable electronic structure of neon 2.8. Valencies, Al = 3 and O = 2 2Al (2.8.3) + 3O (2.6) ==> 2Al3+ (2.8) 3O2– (2.8) can be summarised electronically as 2[2,8,3] + 3[2,6] ==> [2,8] 3+2 [2,8]2–3 so both the aluminium and oxide ions have a full outer shell like a noble gas

TWO

atoms combine with

THREE

atoms to

form Note in this electron diagram, only the original outer electrons are shown above. The three outer electrons of the aluminium atoms (2.8.3) are transferred to the outer shell of the oxygen atom (2.6) until it has a complete octet shell of outer electrons, just like a noble gas 2.8). At the same time, the aluminium ion also attains a stable noble gas electron structure (2.8). Note: The charge on the aluminium ion Al3+ is +3 units (shown as 3+) because there are three more positive protons than there are negative electrons in the aluminium ion. The charge on the oxide ion O2– is –2 units (shown as 2–) because there are two more negative electrons than there are positive protons in the oxide ion.

full electronic structure of aluminium oxide



A GIANT IONIC LATTICE – explaining its properties



The diagram on the right is typical of the giant ionic crystal structure of ionic compounds like sodium chloride and magnesium oxide.



Solid ionic compounds consist of a giant lattice of closely packed ions which are all combine together to form a crystal.



The alternate positive and negative ions in an ionic solid are arranged in an orderly/regular way in a giant ionic lattice structure shown on the right.



The ionic bond is the strong electrical attraction between the oppositely charged positive and negative ions next to each other in the lattice.



The bonding extends throughout the crystal in all directions.



Salts and metal oxides are typical ionic compounds.



This strong bonding force makes the structure hard (if brittle) and have high melting and very high boiling points, so they are not very volatile!



A relatively large amount of energy is needed to melt or boil ionic compounds to reduce/overcome the strong bonding forces. o



Energy changes for the physical changes of state of melting and boiling for a range of differently bonded substances are compared in a section of the Energetics Notes.

The bigger the charges on the ions the stronger the bonding attraction e.g. magnesium oxideMg2+O2– has a much higher melting point than sodium chloride Na+Cl–. o

The ions of magnesium oxide are both doubly charged so the electrostatic attraction is much greater (its actually about 4x as strong attractive force). 

As it happens in this case, the ions in magnesium oxide are smaller than the ions in sodium chloride, so the ions in magnesium oxide can pack closer together and this also increase the attractive bonding force.

o

This double effect results in a much stronger ionic bond in magnesium oxide, so a much greater thermal kinetic energy i.e. a much greater temperature, is required to weaken the giant ionic lattice and melt the crystals of magnesium oxide compared to sodium chloride.

o

Simple experimental evidence – sodium chloride melts at 801 oC, whereas magnesium oxide melts much higher at 2852oC.



Unlike covalent molecules, ALL ionic compounds are crystalline solids at room temperature.



They are hard but brittle, when stressed the bonds are broken along planes of ions

The

crystal structure and properties of Ionic Compounds