A-level Chemistry₇₄₀₅

An element is classified as s, p, d or f block according to its position in the Periodic Table, which is determined by its proton number.

The trends in atomic radius, first ionisation energy and melting point of the elements Na–Ar

The reasons for these trends in terms of the structure of and bonding in the elements.

Students should be able to:

• explain the trends in atomic radius and first ionisation energy
• explain the melting point of the elements in terms of their structure and bonding.

The trends in atomic radius, first ionisation energy and melting point of the elements Mg–Ba

Students should be able to:

• explain the trends in atomic radius and first ionisation energy
• explain the melting point of the elements in terms of their structure and bonding.

The reactions of the elements Mg–Ba with water.

The use of magnesium in the extraction of titanium from TiCl₄

The relative solubilities of the hydroxides of the elements Mg–Ba in water.

Mg(OH)₂ is sparingly soluble.

The use of Mg(OH)₂ in medicine and of Ca(OH)₂ in agriculture.

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

The relative solubilities of the sulfates of the elements Mg–Ba in water.

BaSO₄ is insoluble.

The use of acidified BaCl₂ solution to test for sulfate ions.

The use of BaSO₄ in medicine.

Students should be able to:

• explain why BaCl₂ solution is used to test for sulfate ions and why it is acidified.

AT c and k

PS 2.2

Students could test the reactions of Mg–Ba with water and Mg with steam and record their results.

PS 2.2

Students could test the solubility of Group 2 hydroxides by mixing solutions of soluble Group 2 salts with sodium hydroxide and record their results.

Students could test the solubility of Group 2 sulfates by mixing solutions of soluble Group 2 salts with sulfuric acid and record their results.

Students could test for sulfate ions using acidified barium chloride and record their results.

Research opportunity

Students could investigate the use of BaSO₄ in medicine.

The trends in electronegativity and boiling point of the halogens.

Students should be able to:

• explain the trend in electronegativity
• explain the trend in the boiling point of the elements in terms of their structure and bonding.

The trend in oxidising ability of the halogens down the group, including displacement reactions of halide ions in aqueous solution.

The trend in reducing ability of the halide ions, including the reactions of solid sodium halides with concentrated sulfuric acid.

The use of acidified silver nitrate solution to identify and distinguish between halide ions.

The trend in solubility of the silver halides in ammonia.

Students should be able to explain why:

• silver nitrate solution is used to identify halide ions
• the silver nitrate solution is acidified
• ammonia solution is added.

AT d and k

PS 2.2

Students could carry out test-tube reactions of solutions of the halogens (Cl₂ , Br₂ , I₂ ) with solutions containing their halide ions (eg KCl, KBr, KI).

Students could record observations from reactions of NaCl, NaBr and NaI with concentrated sulfuric acid.

Students could carry out tests for halide ions using acidified silver nitrate, including the use of ammonia to distinguish the silver halides formed.

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

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

Appreciate that society assesses the advantages and disadvantages when deciding if chemicals should be added to water supplies.

The use of chlorine in water treatment.

Appreciate that 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.

Research opportunity

Students could investigate the treatment of drinking water with chlorine.

Students could investigate the addition of sodium fluoride to water supplies.

• cations – Group 2, NH₄ ⁺
• anions – Group 7 (halide ions), OH^– , CO₃ ^2– , SO₄ ^2–

The reactions of Na and Mg with water.

The trends in the reactions of the elements Na, Mg, Al, Si, P and S with oxygen, limited to the formation of Na₂ O, MgO, Al₂ O₃ , SiO₂ , P₄ O₁₀ , SO₂ and SO₃

The trend in the melting point of the highest oxides of the elements Na–S

The reactions of the oxides of the elements Na–S with water, limited to Na₂ O, MgO, Al₂ O₃ , SiO₂ , P₄ O₁₀ , SO₂ and SO₃ , and the pH of the solutions formed.

The structures of the acids and the anions formed when P₄ O₁₀ , SO₂ and SO₃ react with water.

Students should be able to:

• explain the trend in the melting point of the oxides of the elements Na–S in terms of their structure and bonding
• explain the trends in the reactions of the oxides with water in terms of the type of bonding present in each oxide
• write equations for the reactions that occur between the oxides of the elements Na–S and given acids and bases.

AT a, c and k

PS 2.2

Students could carry out reactions of elements with oxygen and test the pH of the resulting oxides.

Transition metal characteristics of elements Ti–Cu arise from an incomplete d sub-level in atoms or ions.

The characteristic properties include:

• complex formation
• formation of coloured ions
• variable oxidation state
• catalytic activity.

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

A complex is a central metal atom or ion surrounded by ligands.

Co-ordination number is number of co-ordinate bonds to the central metal atom or ion.

H₂ O, NH₃ and Cl⁻ can act as monodentate ligands.

The ligands NH₃ and H₂ O are similar in size and are uncharged.

Exchange of the ligands NH₃ and H₂ O occurs without change of co-ordination number (eg Co²⁺ and Cu²⁺ ).

Substitution may be incomplete (eg the formation of [Cu(NH₃ )₄ (H₂ O)₂ ]²⁺ ).

The Cl⁻ ligand is larger than the uncharged ligands NH₃ and H₂ O

Exchange of the ligand H₂ O by Cl^– can involve a change of co-ordination number (eg Co²⁺ , Cu²⁺ and Fe³⁺ ).

Ligands can be bidentate (eg H₂ NCH₂ CH₂ NH₂ and C₂ O₄ ^2– ).

Ligands can be multidentate (eg EDTA^4– ).

Haem is an iron(II) complex with a multidentate ligand.

Oxygen forms a co-ordinate bond to Fe(II) in haemoglobin, enabling oxygen to be transported in the blood.

Carbon monoxide is toxic because it replaces oxygen co-ordinately bonded to Fe(II) in haemoglobin.

Bidentate and multidentate ligands replace monodentate ligands from complexes. This is called the chelate effect.

Students should be able to:

• explain the chelate effect, in terms of the balance between the entropy and enthalpy change in these reactions.

AT d and k

PS 1.2

Students could carry out test-tube reactions of complexes with monodentate, bidentate and multidentate ligands to compare ease of substitution.

PS 2.2

Students could carry out test-tube reactions of solutions of metal aqua ions with ammonia or concentrated hydrochloric acid.

Transition metal ions commonly form octahedral complexes with small ligands (eg H₂ O and NH₃ ).

Octahedral complexes can display cis–trans isomerism (a special case of E–Z isomerism) with monodentate ligands and optical isomerism with bidentate ligands.

Transition metal ions commonly form tetrahedral complexes with larger ligands (eg Cl^– ).

Square planar complexes are also formed and can display cis–trans isomerism.

Cisplatin is the cis isomer.

Ag⁺ forms the linear complex [Ag(NH₃ )₂ ]⁺ as used in Tollens’ reagent.

MS 4.1 and 4.2

Students understand and draw the shape of complex ions.

MS 4.3

Students understand the origin of cis–trans and optical isomerism.

Students draw cis–trans and optical isomers.

Students describe the types of stereoisomerism shown by molecules/complexes.

Transition metal ions can be identified by their colour.

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

d electrons move from the ground state to an excited state when light is absorbed.

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

∆ E = h ν = hc /λ

Changes in oxidation state, co-ordination number and ligand alter ∆ E and this leads to a change in colour.

The absorption of visible light is used in spectroscopy.

A simple colorimeter can be used to determine the concentration of coloured ions in solution.

PS 3.1 and 3.2

Students could determine the concentration of a solution of copper(II) ions by colorimetry.

MS 3.1 and 3.2

Students determine the concentration of a solution from a graph of absorption versus concentration.

AT a, e and k

Students could determine the concentration of a coloured complex ion by colorimetry.

Transition elements show variable oxidation states.

Vanadium species in oxidation states IV, III and II are formed by the reduction of vanadate(V) ions by zinc in acidic solution.

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 reduction of [Ag(NH₃ )₂ ]⁺ (Tollens’ reagent) to metallic silver is used to distinguish between aldehydes and ketones.

The redox titrations of Fe²⁺ and C₂ O₄ ^2– with MnO₄ ^–

Students should be able to:

• perform calculations for these titrations and similar redox reactions.

AT d and k

PS 1.2

Students could reduce vanadate(V) with zinc in acidic solution.

AT b, d and k

PS 4.1

Students could carry out test-tube reactions of Tollens' reagent to distinguish aldehydes and ketones.

AT a, d, e and k

Students could carry out redox titrations.

• the mass of iron in an iron tablet
• the percentage of iron in steel
• the M _r of hydrated ammonium iron(II) sulfate
• the M _r of ethanedioic acid
• the concentration of H₂ O₂ in hair bleach.

Transition metals and their compounds can act as heterogeneous and homogeneous catalysts.

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

The use of a support medium to maximise the surface area of a heterogeneous catalyst and minimise the cost.

V₂ O₅ acts as a heterogeneous catalyst in the Contact process.

Fe is used as a heterogeneous catalyst in the Haber process.

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

A homogeneous catalyst is in the same phase as the reactants.

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

Students should be able to:

• explain the importance of variable oxidation states in catalysis
• explain, with the aid of equations, how V₂ O₅ acts as a catalyst in the Contact process
• explain, with the aid of equations, how Fe²⁺ ions catalyse the reaction between I⁻ and S₂ O₈ ^2–
• explain, with the aid of equations, how Mn²⁺ ions autocatalyse the reaction between C₂ O₄ ^2– and MnO₄ ^–

AT d and k

PS 4.1

Students could investigate Mn²⁺ as the autocatalyst in the reaction between ethanedioic acid and acidified potassium manganate(VII).

In aqueous solution, the following metal-aqua ions are formed:

[M(H₂ O)₆ ]²⁺ , limited to M = Fe and Cu

[M(H₂ O)₆ ]³⁺ , limited to M = Al and Fe

The acidity of [M(H₂ O)₆ ]³⁺ is greater than that of [M(H₂ O)₆ ]²⁺

Some metal hydroxides show amphoteric character by dissolving in both acids and bases (eg hydroxides of Al³⁺ ).

Students should be able to:

• explain, in terms of the charge/size ratio of the metal ion, why the acidity of [M(H₂ O)₆ ]³⁺ is greater than that of [M(H₂ O)₆ ]²⁺
• describe and explain the simple test-tube reactions of: M²⁺ (aq) ions, limited to M = Fe and Cu, and of M³⁺ (aq) ions, limited to M = Al and Fe, with the bases OH^– , NH₃ and CO₃ ^2–

AT d and K

PS 1.2

Students could carry out test-tube reactions of metal-aqua ions with NaOH, NH₃ and Na₂ CO₃

AT d and k

PS 2.2

Students could carry out test-tube reactions to identify the positive and negative ions in this specification.

PS 1.1

Students could identify unknown substances using reagents.

Required practical 11

Carry out simple test-tube reactions to identify transition metal ions in aqueous solution.