AQA A Level Chemistry- Inorganic chemistry

Atomic radius across Period 3

Decreases

First ionisation energy across Period 3

Increases

Exceptions: Drops at Al and S

Melting points across Period 3

Increases from Na to Si

Decreases from Si to Ar

Small rise at S

Explain why the melting point of aluminium is higher than the melting point of sodium

Bigger charge (3+ compared to 1+)

More free/delocalised electrons

Stronger metallic bonding/stronger (electrostatic) attraction between the ions electrons

Explain why the melting point of sulfur is higher than the melting point of phosphorus

S bigger molecule

So more/stronger van der Waals' forces (to be broken or overcome)

Atomic radius down Group 2

Increases

First ionisation energy down Group 2

Decreases

Melting point down Group 2

Decreases

Exception Ca which increases

Explain the melting point of the elements in terms of the structure and bonding of Group 2 elements

Be and Mg have hexagonal close packed structures

Ca, Sr and Ba have cubic structures

The relative solubilities of the hydroxides, X(OH)2, of the elements Mg-Ba in water

Increase down the group

The relative solubilities of the sulfates, XSO4, of the elements Mg-Ba in water

Decrease down the group

The role of magnesium in the extraction of titanium from TiCl4

Mg is used as a reducing agent

Mg is oxidised to 2+ in MgO when it is heated with TiCl4 to about 1200C in an inert atmosphere

Redox equation of Group 2 elements with water

X(s) + 2H2O(l) -> X2+(aq) +2OH-(aq) + H2(g)

Insoluble Group 2 hydroxide

Mg(OH)2

Insoluble Group 2 sulphate

BaSO4

The use of acidified BaCl2 solution

Test for sulfate ions

If Barium Chloride is added to a solution that contains sulphate ions a white precipitate forms (BaSO4)

The use of Mg(OH)2 in medicine

(in solution as milk of magnesia) to neutralise excess acid in the stomach and to treat constipation

The use of BaSO4 in medicine

In a 'Barium meal' given to patients who need x-rays of their intestines. The Barium absorbs the x-rays and so the gut shows up on the x-ray image. Even though Barium compounds are toxic it is safe to use here because of its low solubility

The use of Ca(OH)2 in agriculture

To neutralise acidic soils (slaked lime)

The use of CaO or CaCO3 to remove SO2 from flue gases

Removed by reacting with an alkali such as CaO or CaCO3 slurry (mixed with

water)

• The process is called wet scrubbing:

CaO(s) + 2H2O(l) + SO2(g) -> CaSO3(aq) + 2H2O(l)

CaCO3(s) + 2H2O(l) + SO2(g) -> CaSO3(aq) + 2H2O(l) + CO2(g)

Explain the trend in electronegativity down Group 7

Decreases because the attracting power of the nucleus is shielded by the inner electron shells increasingly down the group

Explain the trend in the boiling point of the elements in terms of their structure and bonding down Group 7

Increases from F to I, because the larger halogens have greater Van der Waals' forces holding the molecules together, as the relative mass increases down the group

Oxidising ability of the halogens down the group

Decreases going down the group and the more reactive halogen will displace a less reactive halogen from a solution of its ions

Reducing ability of the halide ions

Increases as you go down the Group

How to identify and distinguish between halide ions

Reaction with silver nitrate + dilute nitric acid

F- no precipitate

Cl- white precipitate

Br- very pale cream precipitate

I- very pale yellow precipitate

The nitric acid reacts with, and removes, other ions that might also give a confusing precipitate with silver nitrate

Why ammonia solution is added when testing for halide ions

To confirm the precipitate

Solubility of the silver halides in ammonia

AgCl- precipitate dissolves to give a colourless solution

AgBr- precipitate is almost unchanged using dilute ammonia solution, but dissolves in concentrated ammonia solution to give a colourless solution

AgI- precipitate is insoluble in ammonia solution of any concentration

The reaction of chlorine with water to form chloride ions and chlorate(I) ions

Cl2(aq) + H2O(l) -> ClO-(aq) + 2H+(aq) + Cl-(aq)

The reaction of chlorine with water to form chloride ions and oxygen

2Cl2(aq) + 2H2O(l) -> 4HCl + O2

The use of chlorine in water treatment

Cl is pumped into the water in the final stage of water treatment

This water is stored for about 2 hours to allow disinfection to occur

The benefits to health of water treatment by chlorine outweigh its toxic effects

The reaction of chlorine with cold, dilute, aqueous NaOH and uses of the solution formed

Cl2 + 2NaOH -> NaClO + NaCl + H2O

Reaction of sodium and water

2Na + 2H2O --> 2NaOH + H2 , pH>13

Reaction of magnesium and water

Mg + H2O(g) --> MgO +H2 (FAST)

Mg + 2H20(l) --> Mg(OH)2 + H2 , pH~12

Reaction of Na with oxygen

2Na + 1/2O2 --> Na2O(s)

Vigorous

Reaction of Mg with oxygen

Mg + 1/2O2 --> MgO(s)

Vigorous

Reaction of Al with oxygen

2Al + 1.5O2 --> Al2O3(s)

Slow

Reaction of Si with oxygen

Si + O2 --> SiO2(s)

Slow

Reaction of P with oxygen

P4 + 5O2 --> P4O10(s)

Spontaneously combusts

Reaction of S with oxygen (dioxide)

S + O2 --> SO2(g)

Burns steadily

Melting points of Period 3 oxides

Na2O, MgO and Al2O3 have high melting points due to giant ionic lattice structure creating strong forces of attraction between each ion

SiO has a higher melting point than other non-metal oxides due to giant macromolecular structure

P4O1O and SO3 have low melting points due to simple molecular bonding

Reaction of Na2O with water

Na2O + H2O --> 2NaOH, pH 12-14

Reaction of MgO with water

MgO + H2O --> Mg(OH)2, pH 9-10

Reaction of P4O10 with water

P4O10 + 6H2O --> 4H3PO4 (phosphoric(V)acid)

Reaction of SO2 with water

SO2 + H2O --> H2SO3 (sulphuric(IV)acid)

Reaction SO3 with water

SO3 + H2O --> H2SO4 (sulphuric(VI)acid)

Reaction of Al2O3 with water

Insoluble in water but reacts with BOTH acids and bases to form salts - AMPHOTERIC

Reaction of sodium oxide and hydrochloric acid

Na2O + 2HCl --> 2NaCl + H2O

Reaction of magnesium oxide and sulphuric acid

MgO + H2SO4 --> MgSO4 + H2O

Reaction of magnesium oxide and hydrochloric acid

MgO + HCl --> MgCl2 + H2O

Reaction of silicon oxide and sodium hydroxide

SiO2 + 2NaOH --> Na2SiO3 + H2O

Reaction of phosphorus oxide and sodium hydroxide

P4O10 + 12NaOH --> 4Na3PO4 + 6H2O

Reaction of sulphur dioxide and sodium hydroxide

SO2 + 2NaOH --> Na2SO3 + H2O

Reaction of sulphur trioxide and sodium hydroxide

SO3 + 2NaOH --> Na2SO4 + H2O

Reaction of aluminium oxide and hydrochloric acid

Al2O3 + 6HCl --> 2AlCl3 + 3H2O

Reaction of aluminium oxide and sulphuric acid

Al2O3 + 2H2SO4 --> Al2(SO4)3 + 3H2O

Reaction of aluminium oxide and sodium hydroxide

Al2O3 + 2NaOH + 3H2O --> 2NaAL(OH)4

Cause of transition metal characteristics of elements Ti-Cu

An incomplete d sub-level in atoms or ions

The characteristic properties of transition metals Period 4

• complex formation

• formation of coloured ions

• variable oxidation state

• catalytic activity

Ligand

A molecule or ion that forms a co-ordinate bond with a transition metal by donating a pair of electrons

Complex

A central metal atom or ion surrounded by ligands

Co-ordination number

The number of co-ordinate bonds to the central metal atom or ion

Transition metal

A metal that can form one or more stable ions with an incomplete d sub-level

Monodentate ligands

H2O, NH3 (similar in size and are

uncharged) and Cl− (larger than the uncharged ligands NH3 and H2O)

Exchange of the ligands NH3 and H2O occurs without

Change of co-ordination number

Exchange of the ligand H2O by Cl-

Can involve a change of co-ordination number

Types of ligand

Monodentate- one coordinate bond

Bidentate- two coordinate bonds

Multidentate- more than one coordinate bonds

Haem

An iron(II) complex with a multidentate ligand

How oxygen is transported in the blood

Oxygen forms a co-ordinate bond to Fe(II) in haemoglobin

Carbon monoxide is toxic because

It replaces oxygen co-ordinately bonded to Fe(II) in haemoglobin

The chelate effect

Bidentate and multidentate ligands replace monodentate ligands from complexes

An increase in entropy makes the formation of the chelated complex more favourable

Cisplatin complex

Pt(II) with 2 Cl- and 2 ammonia molecules

Square planar

Cis/Z isomer

Tollen's complex Ag+

Diamminesilver(I) complex

Linear complex [Ag(NH3)2]+

Reduction to metallic silver for aldehyde test

Transition metal ions commonly form octahedral complexes with

Small ligands

Transition metal ions commonly form tetrahedral complexes with

Larger ligands (eg Cl-)

Octahedral complexes can display cis-trans isomerism (a special case of E-Z isomerism) with

Monodentate ligands and optical isomerism with bidentate ligands

Why transition metal complexes are coloured

Colour arises when some of the wavelengths of visible light are absorbed and the remaining wavelengths of light are transmitted or reflected

What happens to transition metal ions when light is absorbed

d electrons move from the ground state to an excited state

The energy difference between the ground state and the excited state of the d electrons in transition metals is given by:

∆E = hν = hc/λ

What causes colour change in transition metal ions

Changes in oxidation state, co-ordination number and ligand alter ∆E which leads to the change

Spectroscopy requires use of

Visible light

Colorimetry

Method used to measure the absorption of light by a sample, the more concentrated it is, the more light will be absorbed

Ligand substitution reaction

One ligand is swapped for another

What happens when ammonium vanadate(V) is reduced by zinc in acidic conditions

Yellow to blue as vanadium(V) is reduced to vanadium (IV)

Blue to green as vanadium(IV) is reduced to vanadium(III)

Green to violet as vanadium(III) is reduced to vanadium(II)

How the redox potential for a transition metal ion changing from a higher to a lower oxidation state is influenced

By pH and by the ligand

The redox equation of Fe2+ with MnO4-

MnO4- + 8H+ 5Fe2+ -> Mn2+ +4H2O + 5Fe3

MnO4- half equation

MnO4- +4e- + 8H+ -> Mn2+ +4H2O

Fe2+ half equation

Fe2+ -> Fe3+ + 5e-

The redox equation of C2O42- with MnO4-

MnO4- + 16H+ C2O42- -> 2Mn2+ +8H2O + 10CO2

Heterogeneous catalyst

A catalyst in a different phase from the reactants and the reaction occurs at active sites on the surface

Support medium

Object used to maximise the surface area of a heterogeneous catalyst and minimise the cost

Examples of heterogeneous catalysts

V2O5 in the Contact process.

Fe in the Haber process

Cost implication of heterogeneous catalysts

Can become poisoned by impurities that block the active sites and consequently have reduced efficiency; this has a cost implication

Homogeneous catalyst

Catalyst in the same phase as the reactants

When catalysts and reactants are in the same phase, the reaction proceeds through an intermediate species

How V2O5 acts as a catalyst in the Contact process

Oxidises SO2 to SO3

SO2 + V2O5 -> V2O4 + SO3

Is oxidised back to original state

V2O5 + 1/2O2 -> V2O5

How Fe2+ ions catalyse the reaction between I− and S2O82-

Fe2+ oxidised to Fe3+

2Fe2+(aq) + S2O82-(aq) -> 2Fe3+(aq) + 2SO42-(aq)

Fe3+ oxides iodine to become Fe2+ again

2Fe3+(aq) + 2I-(aq) -> Fe2+(aq) + I2(aq)

How Mn2+ ions autocatalyse the reaction between C2O42- and MnO4-

MnO₄⁻ + 4Mn²⁺ + 8H⁺ → 5Mn³⁺ + 8H₂O + 10CO₂

2Mn³⁺ + C₂O₄²⁻→ 2Mn²⁺+ 2CO₂

Metal -aqua ions limited to M = Fe and Cu

[M(H2O)6]2+

The acidity of [M(H2O)6]3+ is greater than that of [M(H2O)6]2+

Metal -aqua ions limited to M = Al and Fe

[M(H2O)6]3+

The acidity of [M(H2O)6]3+ is greater than that of [M(H2O)6]2+

Why the acidity of [M(H2O)6]3+ is greater than that of [M(H2O)6]2+

Metal 3+ ions are smaller and have a bigger charge than metal 2+ ions. Metal 3+ ions therefore have a higher charge density than metal 2+ ions. This makes the metal 3+ ions much more polarising than the 2+ ions which means the metal 3+ ions can attract electrons from the oxygen atoms ( of the co-ordinatd water molecules) more strongly thus weakening the O-H bond.It is therefore more likely that the hydrogen atom will be released and the more hydrogen ions the more acidic the solution.