Module 3 – Chemistry of the Elements

Includes:

Periodicity- patterns across Period 3 (Na – Ar) Group Chemistry – Groups II , IV , VII

- First Transition Series

Reactions - tests for cations/anions

Periodicity - Blocks in Periodic Table :

Groups I and II are in the s-block Groups III to VII are in the p-block Group 0 (except for He ) is in the p-block 1st Transition series (Sc – Zn) is in the d-block

Periodicity (Na – Ar)

Define: Atomic Radius – definition depends on type of bonding involved:

Covalent Radius – Half the equilibrium distance between two covalently bonded atoms.

Metallic Radius- Half the distance between the nuclei of neighbouring ions in the crystalline metal.

Van der Waals Radius – Half the distance of closest approach between two atoms which are not chemically bonded.

Atomic Radius generally decreases across a period. Across a period the nuclear charge becomes increasingly positive as the number of protons in the nucleus increases. Although the number of electrons also increases, the outer electrons are all in the same shell. The electrons are attracted more strongly to the increasingly positive nucleus, thus reducing the total atomic radius.

Atomic radius generally increases down a group. The outer electrons enter new energy levels down a group, so although the nucleus gains protons, the electrons are not only further away from the nucleus , but are also screened by more inner electron shells. As a result, they are not held so tightly , and the atomic radius increases.

Structures across Period 3

Na,Mg,Al Si P,S,Cl Ar

Metals giant simple discrete atoms

molecules molecular

Electronegativity

This is a measure of the tendency of an atom to attract electrons. It depends on the positive charge in the nucleus. The larger the positive charge, the greater would be the attraction to electrons and hence greater electronegativity.

Electronegativity increases across a period. This is due to increasing positive charge on the nucleus combined with decreasing atomic radius.

Electronegativity decreases down a group. Although the positive charge on the nucleus increases, this is off-set by the increase in atomic radius and the additional screening effects of inner electrons. Fluorine is the most electronegative ,while Caesium is the least.

Melting Points across Period 3

The melting point of a substance is defined as the temperature at which the pure solid is in equilibrium with the pure liquid at atmospheric pressure. Melting point is affected by both the bonding and structure of a substance. When an element melts, the particles break free of the forces holding them together. The greater the force between particles the higher the melting point. The metals Na ,Mg and Al all have metallic structures. The melting points increase from Na to Mg to Al, as attraction between the positive core of ions and the increasing number of delocalised electrons increases. Silicon is a non-metal with the highest melting point in period 3. This is due to the giant molecular structure in which strong

covalent bonds between the atoms hold them tightly in the structure. The non-metals P – Ar have a simple molecular structure and exist as discrete molecules or atoms (Ar). Atoms within the molecules in P – Cl are held together by strong covalent bonds, but the molecules are attracted to each other by weak van der Waals forces. Hence the molecules can be easily separated and these non-metals have low melting points.

Electrical Conductivity across Period 3

This increases from Na to Al. Na has 1 outermost electron which is delocalised, Mg has 2 and Al has 3 . Si is a very poor conductor at room temperature. This is because the four outer electrons are firmly held in four covalent bonds. At higher temperatures more of the outer electrons are promoted to higher energy levels as they gain energy and become delocalised. Hence , Si is a semi-conductor. The other elements (P – Ar) do not conduct electricity as their electrons are held in covalent bonds and are not free to move.

Density across Period 3

Density is related to the size of the atoms and their packing.

Density = mass/volume ; volume = 4/3 πr3 (r= atomic radius)

For the metals, the effects of atomic radius are more important than packing, which is similar in metals. Across a period there is a decrease in atomic radius and so more atoms can be found in a given volume. Hence the density of metals increases across the period. The transition metals have the highest density. Across a period the density reaches a maximum in the middle of the table. The maximum in period 2 is carbon which is strongly covalently bonded in diamond. The maximum in period 3 is Aluminium. To the right of this maximum the density of the elements falls as they exist as small covalent molecules which are widely spread with weak intermolecular forces.

Diagonal Relationships

Periodic properties show that a trend in atomic properties going left to right e.g. Atomic radius will tend to be cancelled out by moving down the group. Hence elements along diagonal lines from top left to bottom right will have similar atomic sizes and reactivities.

Reactions of elements of period 3 with Oxygen

Sodium

Sodium burns in oxygen with an orange flame to produce a white solid mixture of sodium oxide and sodium peroxide.

For the simple oxide:

For the peroxide:

Magnesium

Mg is normally covered with a layer of its oxide. It burns rapidly in air or oxygen with a brilliant white flame. This reaction is used in fireworks and flares.

N.B. If Magnesium is burnt in air rather than pure oxygen, it also reacts with the nitrogen in the air a mixture of magnesium oxide and magnesium nitride is obtained. (3Mg + N2 → Mg3N2)

Aluminium

Al metal is normally coated with a layer of its oxide which forms very rapidly when exposed to air. The oxide layer is tough and protects the underlying metal from corrosion.

Silicon

Si burns in oxygen to form a white solid, silicon dioxide , which is called silica and has a giant covalent lattice.

Phosphorus

Phosphorus reacts very rapidly with oxygen and would ignite without any external heating when exposed to air. In a limited supply of oxygen, the main product is the white solid phosphorus (III) oxide:

With plenty oxygen the main product is the white solid phosphorus (V) oxide:

Sulphur

Sulphur burns easily in air or oxygen with a blue flame to form the gas sulphur dioxide:

N.B. Sulphur dioxide reacts with oxygen when passed over a heated platinum catalyst to give Sulphur (III) oxide:

2SO2 + O2 2SO3

Chlorine does form oxides but not by direct reaction with oxygen.

Argon is extrememly unreactive and has not yet been persuaded to form any oxides.

Reactions of elements of Period 3 with Chlorine

Sodium

Sodium burns in chlorine with a bright orange flame. White solid sodium chloride is produced.

Magnesium

Magnesium burns with its usual intense white flame to give white magnesium chloride.

MgO and MgCl2, as well as NaCl and Na2O have giant ionic structures.

Aluminium

Al reacts rapidly when heated in chlorine to give the white solid Aluminium Chloride:

This is an easily vaporised solid (sublimes) which fumes in moist air and is soluble in organic solvents. Hence it has a covalent character. In the vapour phase it consists of dimeric molecules of Al2Cl6.

When heated:

Al2Cl6(g) 2AlCl3(g)

Silicon

If chlorine is passed over silicon powder heated in a tube, it reacts to produce silicon tetrachloride. This is a colourless liquid which vaporises and can be condensed further along the apparatus. It is covalently bonded.

Phosphorus

White phosphorus burns in chlorine to produce a mixture of two chlorides, phosphorus(III) chloride and phosphorus(V) chloride (phosphorus trichloride and phosphorus pentachloride).

Phosphorus(III) chloride is a colourless fuming liquid.

Phosphorus(V) chloride is an off-white (going towards yellow) solid.

Sulphur

If a stream of chlorine is passed over some heated sulphur, it reacts to form an orange, evil-smelling liquid, disulphur dichloride, S2Cl2.

Chlorine – no reaction.

Argon- no reaction.

Reactions of elements of Period 3 with water

Sodium

Sodium melts, fizzes and floats in a rapid exothermic reaction with cold water to give Hydrogen gas and Sodium Hydroxide: 2Na(s) + 2H2O(l) → 2NaOH (aq) + H2(g)

The solution is highly alkaline with a pH of 12-14 due to the high concentration of OH- ions present.

Magnesium

Magnesium has a very slight reaction with cold water, but burns in steam.

A very clean coil of magnesium dropped into cold water eventually gets covered in small bubbles of hydrogen which float it to the surface. Magnesium hydroxide is formed as a very thin layer on the magnesium and this tends to stop the reaction.

The solution is weakly alkaline with a pH of 9-11 since magnesium hydroxide is not as soluble as NaOH.

Magnesium burns rapidly in steam with its typical white flame to produce white magnesium oxide and hydrogen.

Both reactions are redox reactions.

Aluminium

Al does not react readily with water or even steam at high temperatures. The outer layer consists of a thin strongly bonded layer of Al2O3 which protects the metal underneath from attack by water. Hence, aluminium is used as a strong, light metal in ships, aircraft, and light cooking vessels.

Silicon – no reaction

Phosphorus -no reaction

Sulphur – no reaction

Chlorine

Chlorine reacts slowly with water to form HCl and chloric(I) acid which readily decomposes to hydrochloric acid and oxygen.

Cl2(g) + H2O(l) → HCl(aq) + HClO (aq)

Chloric(I) acid

2HClO (aq) → 2HCl(aq) + O2(g)

Argon – no reaction

Oxidation numbers of chlorides and oxides across Period 3

Formulae:

Na2O MgO Al2O3 SiO2 P4O6 SO2 Cl2O

Na2O2 MgCl2 AlCl3 SiCl4 P4O10 SO3 ClO2

NaCl PCl3 S2Cl2 Cl2

PCl5

In these compounds the oxygen and chlorine are more electronegative than the other elements. This means that across the period the elements have positive oxidation numbers in their oxides and chlorides. An account of oxidation number by group is as follows:

Group 1 : Na forms only the ion Na+ with oxidation no. +1 Group 2: Mg forms only the ion Mg2+ with oxidation no. +2 Group 3: Al forms the ion Al3+ or shares electrons and has oxidation no. +3 Group 4: Si shares electrons and has oxidation no. +4 Group 5: P has several oxidation nos. with the highest at +5. In P4O6 it is +3.

In P4O10 and PCl5 it is +5 Group 6 : S has several oxidation nos. with the highest at +6. In S2Cl2 it is +1.

In SO2 it is +4. In SO3 it is +6. Group 7 : Cl has several oxidation nos. with the highest at +7. In Cl2O it is

+1. In ClO2 it is +4. In Cl2O7 it is +7.

Reactions of Chlorides across Period 3 with water

NaCl dissolves in water to produce a neutral solution while MgCl2 produces a slightly acidic solution. AlCl3 produces a solution which may be very acidic with a pH as low as 3. The increasing acidity of solutions of chlorides from sodium to aluminium is explained by the decreasing size and increasing charge of the positive ions present. The small, highly charged ions e.g. Al3+ strongly polarise the OH bonds in the surrounding water molecules. This causes some of the water molecules to give up H+ ions producing an acidic solution.

The chlorides of Silicon, Phosphorus and Sulphur all react with water to form acidic products. Both Silicon and Phosphorus chlorides react to form hydrochloric acid.

SiCl4(l) + 4H2O(l)→ Si(OH)4(aq) + 4HCl(aq)

PCl3(l) + 3H2O(l) → H3PO3(aq) + 3HCl(aq)

The hydrolysis of the chlorides of sulphur is more complex, forming sulphur, hydrogen sulphide and sulphite ions and hydrochloric acid.

Period 3 Chlorides

Formula NaCl MgCl2 AlCl3 SiCl4 PCl3 S2Cl2 Cl2

PCl5

State at 20⁰ C

solid solid solid liquid liquid liquid gas

Electrical Conductivity

good good poor none none none none

Structure Giant ionic

Giant ionic

Simple covalent

Simple covalent

Simple covalent

Simple covalent

Simple covalent

Reaction with water

soluble

soluble

HCl (g)

produced

HCl(g)

produced

HCl(g)

produced

HCl(g)

produced

Slight

reaction

pH of soln. neutral

Slightly acidic

acidic acidic acidic acidic acidic

Reactions of oxides across Period 3 with water

The acid-base properties of the oxides are linked to their structures. At the left-hand end of the period the oxides of the electropositive elements are highly basic and form alkaline solutions. Oxide ions remove H+ ions from water. Hence, sodium oxide (Na2O) reacts with water to form sodium hydroxide.

Na2O(s) + H2O(l) →2 NaOH(aq)

MgO has a slight reaction to form magnesium hydroxide.

MgO(s) + H2O(l) → Mg(OH)2(aq)

Al2O3 is amphoteric. It reacts with acidic solutions to form Al3+ salts and with alkali to form aluminate.

Al2O3(s) + 6H+ (aq) → 2Al3+(aq) + 3H2O(l) : basic character

Al2O3(s) + 2OH-(aq) + 3H2O(l) → 2Al(OH)4- (aq) : acidic character

SiO2 behaves as a very weak acid. It does not react with water. However, it reacts with a base like NaOH forming the silicate ion.

SiO2(s) + 2OH-(aq) → SiO32- (aq) + H2O(l)

Silicate

Oxides of Phosphorus sulphur and chlorine at the right-hand side of the period show acidic behaviour.

P4O10(g) + 6H2O(l) → 4H3PO4(aq) (phosphoric(V)acid

SO2(g) + H2O(l) → H2SO3(aq) sulphurous acid

SO3(g) + H2O(l) → H2SO4(aq) sulphuric acid

Cl2O7(l) + H2O(l)→ 2HClO4(aq) chloric(I) acid

Period 3 Oxides

Formula Na2O MgO Al2O3 SiO2 P4O10

(P4O6)

SO2

(SO3)

Cl2O7

(Cl2O)State at 20⁰C

solid solid solid solid Solid

(solid)

Gas

(gas)

Liquid

(gas)Elect.

Cond. Liq.

state

good good good Very poor

none none none

Reaction with water

Forms NaOH(aq)

Forms Mg(OH)2

Does not react

Does not react

Forms H3PO4

Forms H2SO3

(H2SO4)

Forms HClO4

Nature of oxide

basic Weakly basic

amphoteric

acidic acidic acidic acidic

N.B. Al2O3 and SiO2 do not dissolve in water.

Group Chemistry

Group II – Alkaline Earth Elements

Element Electronic configuration

Atomic radius/pm

m.p./⁰C 1st I.E.(kJ/mol)

2nd I.E.(kJ/mol)

Ionic Radius/pm

S.E.P./V

Be 1s22s2 111 1278 900 1760 31 -1.75Mg [Ne]3s2 160 651 736 1460 65 -2.37Ca [Ar]4s2 197 850 590 1150 99 -2.87Sr [Kr]5s2 215 770 548 1060 113 -2.89Ba [Xe]6s2 217 704 502 966 135 -2.91

Just know the electronic configurations up till Ca!!

General Properties

1. All have 2 electrons in their outermost shell and 8 electrons in their penultimate shell except Be which has 2 electrons. They all form a stable M2+ cation with a noble gas structure. Their reactivity increases with atomic number. This is due to the increasing ease with which electrons can be lost from the atoms. Be is unusual. It does not easily lose 2 electrons to form Be2+ and many compounds of Be are covalent. Be2+ is small and highly charged and hence has a high charge density which explains its tendency to form covalent bonds and the amphoteric nature of its oxide.

2. Ionisation energy data show that while it is slightly more difficult to remove the second electron in the outer s-subshell than the first, the main differences come from the removal of the third electron from the penultimate shell. Ionisation energy gets smaller down the group as the no. of shells increases and the attractive force of the nucleus is felt less strongly by the outer electrons which are screened more effectively from the positive charge.

3. Elements of Groups 1 and 2 have larger atomic radii than other elements in their periods. The outer s-electrons are held relatively weakly by the nucleus. Because of the increase in atomic size, the attractive forces between atoms in the metal lattices are also weakened and so m.p. and b.p. are relatively low for metals.

4. The Ionisation energy data and electrode potential values both show that the reactivity of Group 2 elements increases with atomic number.

Reactions of elements with water

The rate of reaction increases from Be to Ba for two reasons:

1. From Be to Ba the metals show a greater negative Standard Electrode Potential.

2. Increased solubility of the hydroxides.

Be does not react at all with water ,cold or hot, because of the insolubility of its hydroxide. Mg reacts slowly with cold water. With hot water it slowly forms MgO and H2(g).

Mg(s) + 2H2O(l) → Mg(OH)2(aq) + H2(g)

When steam is passed over heated magnesium , an exothermic reaction occurs to give MgO and H2(g).

Mg(s) + H2O(g) → MgO(s) + H2(g)

Calcium is more reactive than Mg. It reduces cold water to form an alkaline solution of calcium hydroxide known as Lime Water.

Ca(s) + 2H2O(l) → Ca(OH)2(aq) + H2(g)

Strontium reacts more quickly than calcium.

Barium reacts vigorously but less than sodium.

M(s) + H2O(l)→ M(OH)2(aq) + H2(g)

M = Ca,Sr,Ba.

Reactions of elements with air or oxygen

All group 2 elements tarnish in air forming a coat of oxide with the speed of reaction increasing from Be to Ba. The coat or film is protective for Be and Mg but not the others which continue reaction and absorb water and carbon dioxide to eventually form carbonates:

O2 H2O CO2

M MO M(OH)2 MCO3

Reaction stops at MO for Be and Mg.

Reaction continues to M(OH)2 and then MCO3 for Ca,Sr,Ba.

Ca,Sr, and Ba are stored under paraffin oil. When all group 2 metals are heated in air or oxygen they burn vigorously to form a white ionic oxide.

e.g. 2 Mg(s) + O2(g) → 2MgO(s)

Reactions of elements with acids

The group 2 elements generally react with dilute acids to release hydrogen gas. The reactivity increases down the group. The reactions are redox reactions.

Oxid. No. 0 +1 +2 0 Mg(s) + 2H+(aq) → M2+(aq) + H2(g)

N.B. In their reactions with dilute sulphuric acid Ca , Sr and Ba react with decreasing vigour. This is due to the decreasing solubility of the sulphate MSO4 , which retards the reaction.

Reactions of Group 2 Oxides with water

The oxides and hydroxides of group 2 elements are basic and neutralized by acids.

O2- (s) + 2H+(aq) → H2O(l)

OH- (aq) + H+ (aq) → H2O(l)

The solubility of the hydroxides in water increases down the group. Thus Mg(OH)2

is insoluble in water and Ba(OH)2 is soluble. BeO neither reacts with nor dissolves in water. The others react forming hydroxides, MgO slowly and CaO, SrO, BaO with great vigour and evolution of much heat increasing down the group. For example Quicklime (CaO) reacts with water to form Ca(OH)2 or lime water , which is sparingly soluble in water.

CaO(s) + H2O(l) → Ca(OH)2(aq) Quicklime Lime Water

When water is dropped slowly onto calcium oxide the temperature of the reaction may be high enough to boil any excess water. This process is called slaking .

Uses of Oxides and Hydroxides

1. Solid calcium hydroxide is used to neutralize acidic salts e.g. to reduce acidity in some lakes.

2. A suspension of Mg(OH)2 is called milk of magnesia and is used to cure indigestion.

3. Lime-water – used as a test for CO2 .

Ca(OH)2 (aq) + CO2(g) → CaCO3(s) + H2O(l) Lime-water white ppt.

When excess CO2(g) is bubbled through lime-water , the ppt. dissolves to give Calcium hydrogen carbonate.

CaCO3(s) + H2O(l) + CO2(g) → Ca(HCO3)2(aq)

Thermal Stability of Carbonates and Nitrates

Carbonates:

The thermal decomposition and stability of ionic compounds depends on the forces that exist between positive and negative ions in the lattice. The strength of this force of attraction is estimated from the lattice energy. The lattice energy is always negative and the more negative it is , the greater is the attraction between ions. It depends on three factors:

a. The greater the charge on the ion the greater is the attraction for an ion of opposite charge and the greater is the L.E.

b. The smaller the ions are the more closely they approach each other in the lattice and the more negative is the lattice energy and hence larger melting points.

c. The smaller the cations the greater is its polarizing power and hence ability to distort larger anions and break them up giving lower melting points.

Group II Carbonates

Decomposition temp/⁰c

MgCO3 Value of 400 ThermalCaCO3 L.E. becomes 900 stabilitySrCO3 more 1280 increasesBaCO3 negative 1360 heate.g. MgCO3(s) → MgO(s) + CO2(g)

The thermal stability of the carbonates increases down the group for two reasons:

1. Polarisation: Small highly charged positive ions e.g. Mg2+ have the greatest polarizing power and distort large negative anions e.g. CO3

2- by pulling the oxygen atoms towards them. Once the carbonate is highly polarized it forms an oxide and CO2. Ba2+ has the least tendency to distort CO3

2- ions.

2. Lattice energy of the oxide formed: The lattice energy is more exothermic with a small cation than with a large cation. In the decomposition of MgCO3 the formation of MgO with a high lattice energy can be considered as the driving force for the reaction. MgO has a high lattice energy. It is very stable and is favoured to be formed. BaCO3 has CO3

2- not highly distorted and the product BaO has a lower lattice energy.

Thus, the order of decreasing L.E. is –

MgO > CaO > SrO > BaO

Nitrates:

Nitrates decompose on heating in a Bunsen flame. Group 1 nitrates, except Lithium Nitrate, form the corresponding Nitrites which are then stable to heat.

e.g. 2KNO3 (s) → 2KNO2 (s) + O2(g)

The small decrease in size from NO3- to NO2

- is enough to achieve thermal stability with Group 1 cations, except Lithium. In contrast, Lithium Nitrate and all Group 2 Nitrates decompose on heating to form the oxide, NO2 and O2 .

e.g. 2Mg(NO3)2 (g) → 2MgO(s) + 4NO2(g) + O2(g)

4LiNO3 (s) → 2Li2O(s) + 4NO2(g) + O2(g)

LiNO3 reacts like Group 2 Nitrates. This is an example of the Diagonal relationship.

Group 2 cations and Li+ need the much smaller O2- ion to achieve thermal

stability. O2- is smaller and more highly charged than NO3

- and the oxide formed has a high lattice energy. The stability of Group 2 Nitrates increases down the group. This is due to a decreased polarizing power of the cation as the ionic radius increases.

Solubility of Group 2 Sulphates:

Solubility is given by ∆Hsoln. (enthalpy of solution). The more negative the ∆Hsoln , the more soluble is the compound.

∆Hsoln is given by: ∆Hsoln = -L.E. + ∆Hhydration

Two factors affect solubility:

1. –L.E. is highly endothermic2. ∆Hhydration (both ions) is highly exothermic.

Decrease in ∆Hhydration of cation

L.E. do not vary much, since in size anions are much bigger than cations.

The solubility of the sulphates decreases down the group. The enthalpy of hydration of SO4

2- is the same for all. The lattice energy of the Group 2 sulphates becomes less exothermic as the size of the metal increases but the change is small. The change in Lattice Energy from MgSO4 to BaSO4 is small because of the large size of the SO4

2- ions . The difference in ∆Hhydration of the gaseous cation from Mg2+ to Ba2+ is very large because of the large difference in ionic radius from the small Mg2+ to the large Ba2+. It is the ∆Hhydration of the cations that determines the solubility trend. Thus the enthalpy of solution (∆Hsoln ) becomes less exothermic down the group and hence solubility decreases.

Uses of Group 2 Compounds

Uses of the oxides: BeO is used to make porcelain crucibles because of its refractory (ability to take heat) nature i.e. high melting point and low thermal expansion. It is also an excellent electrical insulator and is used to make spark plugs. MgO and CaO are used to male refractory bricks to line the inside of furnaces. CaO is also used to make Ca(OH)2 .

Uses of the hydroxides: A suspension of Mg(OH)2 in water known as Milk of Magnesia, is used as an antacid. Ca(OH)2 is used to make mortar. It is also used to make bleaching powder and NaOH. It is also used for neutralizing acids in lakes , and in water softening.

Group IV

General Properties:

Element C Si Ge Sn Pbm.p./°C 3652(graphite) 1410 957 232 325Electron Config.

[He]2s22p2 [Ne]3s23p2 [Ar]3d104s24p2 [Kr]4d105s25p2 [Xe]5d106s26p2

structure Giant molecular

Giant molecular

Giant molecular

Giant metallic Giant metallic

Electrical conductivity

Graphite(good)

Diamond(none)

semi semi good good

Principal oxidation no.

+4 +4 +2,+4 +2,+4 +2,+4

Group IV is very diverse. C and Si are non-metals, Ge is a semi-metal, Sn and Pb are metals. The +2 oxidation state becomes more stable down the group.

Physical Properties

The bonding between atoms of Group 4 elements changes from covalent in the Giant Molecular C and Si to Giant Metallic in Sn and Pb. Giant molecular structures have high m.p. Giant metallic structures conduct heat and electricity very well. As the atoms get larger down the group, the bonding between atoms gets weaker and attraction of nuclei for electrons gets weaker. The metallic character increases down the group. With a larger atomic radius the outer electrons are further from the nucleus and better shielded by inner shells of electrons. Hence, Sn and Pb have structures with positive Pb and Sn ions in a sea of delocalized electrons.

Inert Pair Effect

Carbon, C 1s22s22p2 (2.4) [He]2s 2p

Silicon, Si 1s22s22p63s23p2 (2.8.4) [Ne]3s 3p

Germanium, Ge [Ar]3d104s24p2[Ar]3d 4s 4p

Tin, Sn [Kr]4d105s25p2[Kr]4d 5s 5p

Pb [Xe]5d106s26p2 [Xe]5d 6s 6p

Down the group there is an increasing tendency for the outer pair of s electrons to remain inert , so that they do not take part in bonding and the +2 oxidation state becomes more stable. In Sn and Pb the outer electrons are held more weakly due to screening effects of the inner shells of electrons. The outer s-electrons penetrate to some extent the shell immediately below them. Thus, Tin and Lead tend to lose 2 electrons in ionic bonding. In the Pb2+ ion only the 2p electrons are lost. The s-electrons remain as a part of the inner core and is called an Inert Pair.

E⁰ values show this as well:

Sn4+(aq) + 2e → Sn2+(aq) E⁰ = +0.15 V

Pb4+(aq) + 2e → Pb2+(aq) E⁰ = +1.69V

The values show the greater stability of the +2 oxidation states of Sn and especially Pb.

Trends in Melting Point

The very high melting point of diamond is caused by the great energy needed to break the strong C-C covalent bonds. From C to Ge, bond lengths increase and bond strengths decrease, hence Silicon and Ge have lower m.p. There is a further fall to Sn and Pb. Sn and Pb have metallic structures which are weaker than the strong covalent bonds. Pb has a slightly stronger metallic bond than Sn and hence its m.p. is slightly higher.

Tetrachlorides of Group 4

Formula CCl4 SiCl4 GeCl4 SnCl4 PbCl4

Structure simple molecular (tetrahedral)

Thermal Stability stable to high temps. Decomposes Decomposes

on heating at room temp.

SnCl4→SnCl2+Cl2 PbCl4→PbCl2+Cl2

Reaction with water No rxn. Hydrolysed readily to form hydroxides and HCl(g)

Bonding

The chlorides are all simple molecular with tetrahedral shapes. The stability of the chlorides decreases down the group and the +2 oxidation state becomes more stable. Only Tin and Lead form chlorides in which oxidation state is +2. The other chlorides exist in the +4 state only. Tin(II) Chloride is a solid that is soluble in water , giving a solution which conducts electricity. It is also soluble in organic solvents and its melting point is 246⁰C. Lead (II) Chloride is a solid, sparingly soluble in water and gives a solution which conducts electricity. It melts at 501⁰C. Hence, Tin(II)Chloride has both covalent and ionic character, and Lead(II)Chloride is mainly ionic.

Thermal Stability of the Tetrachlorides

Thermal stability decreases down the group. This is explained in terms of the bond strengths and bond lengths of the X-Cl bonds. A covalent bond is formed by overlap of atomic orbitals. In the Pb-Cl bond there is little overlap because of differences in energy between the orbitals. Hence, this covalent bond is very weak and is easy to break. The Pb-Cl bond is much longer than the other X-Cl bonds.

Hydrolysis of Tetrachlorides (Reaction with water)

XCl4(l) + 2H2O(l) → XO2(s) + 4HCl(g)

X = Si, Ge, Sn, Pb.

Except for Carbon Tetrachloride (CCl4) ,all the other tetrachlorides are readily hydrolysed to form the dioxide as a precipitate and HCl(g). The tetrachlorides are all liquids at room temperature.

Why does CCl4 not react with water?

Cl Cl

Cl X∂+ Cl Cl X∂+ Cl

Cl ∂- O H Cl O H

H H

Unstable Intermediate

Cl

Cl X OH + HCl(g)

Cl

Reaction continues three more times to give X(OH)4 :

XCl4(l) + 4H2O(l) → X(OH)4 (aq + 4HCl(g)

Followed by:

X(OH)4 (aq) → XO2(s) + 2H2O(l)

Overall: XCl4(l) + 2H2O(l) → XO2(s) + 4HCl(g)

Explanation: In hydrolysis a lone pair of electrons from the oxygen atom of a water molecule attacks the Group 4 atom in the chloride which has a partial positive charge due to the electron-withdrawing effect of the four chlorine atoms. The Group 4 atom briefly has 5 electron pairs around it. For this to occur there must be empty orbitals present. In silicon and the lower elements, the extra pair of electrons is accommodated using the available empty d-orbitals. Due to its

electronic configuration (1s22s22p2) Carbon has no available d-orbitals and hence cannot form the intermediate and so no hydrolysis is possible with CCl4.

Oxides with the +4 oxidation state

Formula CO2 SiO2 GeO2 SnO2 PbO2

Structure simple giant giant molecular/ionic

Molecular molecular

Thermal stable to high temperatures decomposes on

Stability heating:

PbO2→PbO +1/2 O2

Reactions XO2 + 2OH- → XO32- + H2O

With bases e.g. SiO2 + 2OH- → SiO32- + H2O

Dil.alkali conc alkali increasingly severe conditions fused alkali

Reactions no rxn. No rxn. XO2 + 4HCl → XCl4 + 2H2O

With acids (X = Ge,Sn,Pb)

Oxidizing very weak increase in oxidizing power powerful

Power

Points to note:

1. All Group 4 elements form a dioxide with empirical formula XO2 . There is a distinct change in the chemical properties of these oxides down the group. This is due to a change in the stability of the +4 oxidation state. Down the group the +2 state becomes more stable.

2. The oxides show a trend in structure from molecules of CO2 to giant structures intermediate between ionic and covalent down the group.

3. The thermal stability decreases down the group. Strong heating of PbO2

gives O2 and PbO. The +2 state is more stable for lead. SiO2 is very thermally

stable. It is often a component of ceramics. Ceramics can withstand high temperatures but are brittle. One crystalline form of SiO2 is quartz which is a hard , brittle ,clear , colourless solid.

4. Since Group 4 elements form two sets of oxides, XO2 and XO, it is possible to convert XO2 to XO , which is a reduction reaction. Hence the dioxides,XO2

are oxidizing agents. PbO2 and SnO2 are the most powerful of these oxidizing agents and both can oxidize HCl to form chlorine.e.g. PbO2(s) + 4HCl(aq) → PbCl2(s) + Cl2(g) + 2H2O(l)

5. Acid-Base Reactions: Covalent oxides are often acidic and ionic oxides are basic. The metal oxides are often basic and non-metal oxides are often acidic. Hence the oxides at the top of the group CO2 and SiO2 have an acidic nature e.g. CO3

2- is produced in dilute aqueous alkaline solutions. CO2 (g) + 2OH-(aq) → CO3

2- (aq) + H2O(l)

The ease of formation of the XO32- ion decreases down the group as the

acidic character decreases and the ionic character increases. The oxides of Ge, Sn and Pb have an increasing degree of ionic character and hence show basic properties. Hence these oxides are amphoteric and form salts with acids. PbO2 shows some acidic character when it reacts with conc. NaOH to form the plumbate(IV) ions.

PBO2(s) + 2NaOH(aq) → Na2PBO2(aq) + H2O(l)

Sodium plumbate(IV)

Oxides in the +2 oxidation state

Formula CO GeO SnO PbO

Structure simple molecular giant ionic

Acid-base neutral amphoteric

Behavior e.g. SnO + 2H+ (aq) → Sn2+(aq) + H2O(l)

SnO + 2OH-(aq) → SnO22-(aq) + H2O(l)

Stannate(II)

increasing basic character

increasing ionic character

N.B. SiO does not exist

Points to note:

1. Carbon monoxide has completely different properties to those of the other oxides with a +2 oxidation state. It is a covalently bonded gaseous compound and is a neutral compound. The other three are ionically bonded solid oxides that show increasingly basic properties as the ionic character of the oxides increases.

2. Oxidation of the monoxides: Since PbO2 is the least thermally stable of the dioxides, Pb forms the most stable monoxide PbO. The monoxides CO,GeO, and SnO all react readily with oxygen in air to give the respective dioxides. 2CO(g) + O2(g) → 2CO2(g) 2GeO(s) + O2(g) →2GeO2(s)PbO will not form PbO2 when heated in air, but at 400⁰C with prolonged heating it forms Red Lead (Pb3O4). Red Lead, Pb3O4 is dilead(II)lead(IV)oxide i.e. 2 moles of lead(II)oxide = PbO PbO 1 mole of lead(IV)oxide = PbO2

Total = Pb3O4

3. Acid-Base reactions of monoxides: The oxides become more basic down the group. This is because the bonding in the oxides become more ionic down

the group and ionic oxides show basic properties. Even though the basic character increases PbO is not a basic oxide but is amphoteric. In acids PbO forms Pb2+ salts, with alkali it forms plumbate(II) : PbO(s) + 2HNO3(aq) → Pb(NO3)2(aq) + H2O(l) PbO(s) + 2NaOH(aq)→ Na2PbO2(aq) + H2O(l) Sodium plumbate(II)

Reducing Properties

Sn4+ (aq) + 2e ⇄ Sn2+(aq) +0.15VPb4+(aq) + 2e ⇄ Pb2+(aq) +01.64VGeCl2 and SnCl2 are both reducing agents. Acidified aqueous SnCl2 reduces Fe3+ ions and dichromate(VI) ions (Cr2O7

2-). Down the group there is an increase in the stability of the +2 state as reflected in the E° values. The highly positive E° value for the reduction of Pb4+ indicates that lead(IV) is very easily reduced to Pb2+ and hence lead(IV) is an oxidising agent.

Silicon Dioxide – a Giant Molecular Solid

Formula =( SiO2)n - silica, sand ,quartz.

Unit Cell:

Each Silicon ato mis covalently bonded to four oxygen atoms, similar to the diamond structure. Like diamond, SiO2 is a hard solid, insoluble in wáter and a non-conductor of electricity. The Si-O bond has a high bond energy and compounds of Silicon containing other bonds have a strong tendency to be converted into compounds containing SiO2. Silicon differs from carbón as it does not form long chains of silicon atoms, since the Si-Si bond is much less stable tan the Si-O bond.