Inorganic Chemistry Practicals (AQA A Level Chemistry): Flashcards

False.

Sulfur initially forms SO₂. Further oxidation to SO₃ is slow unless a vanadium catalyst is used, so sulfur does not always reach its highest oxidation state under typical laboratory conditions.

The pH decreases (becomes more acidic) across the period. Sodium oxide dissolves to give a strongly alkaline solution (pH _{13–14). Aluminium oxide is amphoteric (}pH 7). Silicon, phosphorus, and sulfur oxides produce progressively more acidic solutions.

In the Period 3 oxides practical, elements such as Na, Mg, and S are lowered into oxygen using a deflagrating spoon inside a combustion jar. The resulting oxides are added to water and the pH is tested.

False.

White phosphorus spontaneously catches fire in air (spontaneous ignition). Red phosphorus also reacts but less vigorously.

A pH probe (or pH meter) and pH paper (or universal indicator paper) can both be used to measure or estimate the pH of the resulting solution.

Chelate effect is the greater stability of complexes containing bidentate or multidentate ligands compared to those with monodentate ligands, arising from the entropy increase when each multidentate ligand displaces more than one water molecule.

When excess aqueous ammonia is added to [Cu(H₂O)₆]²⁺, the solution changes from pale blue to deep blue. This colour change reflects a change in both ligand and coordination environment.

False.

Both the colour and whether a solution or precipitate forms must be recorded. This information together is necessary to correctly identify the complex or product.

Each multidentate ligand displaces more than one water molecule, increasing the number of free particles in solution. This increases entropy, making complex formation more thermodynamically favourable.

In a UV/visible spectrophotometry experiment to compare rates of ligand substitution, absorbance is proportional to concentration, so as the complex concentration changes during the reaction, the absorbance reading changes proportionally.

False.

Concentrated HCl causes ligand substitution: Cl^- ligands replace water, forming [CuCl₄]^2-, which is yellow-green. The green appearance is a mix of the blue aqua ion and yellow-green tetrachlorocuprate.

To stop (quench) the reaction by greatly diluting the concentrations of all species. This allows the absorbance at that moment in time to be measured without the reaction continuing in the cuvette.

Sulfuric acid does not react with or oxidise the manganate(VII) ions under titration conditions. Hydrochloric acid would be oxidised to Cl₂ by MnO₄^-, and nitric acid is itself an oxidising agent that would interfere with the titration.

In a KMnO₄ titration, the endpoint is seen as a pale pink tinge in the flask. This appears because MnO₄^- is purple and acts as its own indicator, requiring no separate indicator.

True.

From the balanced equation: MnO₄^- + 8H⁺ + 5Fe²⁺ → Mn²⁺ + 4H₂O + 5Fe³⁺. One mole of manganate(VII) reacts with five moles of iron(II).

Moles of Cr₂O₇^2- = (0.0180 × 28.50) / 1000 = 5.13 × 10⁻⁴ mol

Moles of Fe²⁺ = 6 × 5.13 × 10⁻⁴ = 3.08 × 10⁻³ mol

In a K₂Cr₂O₇ titration with Fe²⁺, the colour change at the endpoint is from orange to bluish green. The indicator sodium diphenylaminesulfonate is often used to give a clearer colour change (colourless to purple).

False.

A burette with white numbering should be used. The deep purple of KMnO₄ makes it very difficult to read black numbers against the coloured solution.

Mn²⁺ (aq) is very pale pink, effectively colourless at typical concentrations. The pale pink tinge at the endpoint is due to a slight excess of MnO₄^- after all Fe²⁺ has been oxidised.

A titration in which the reaction between oxidising and reducing agents is used to determine the concentration of an unknown solution, with the endpoint identified by a colour change in the solution or an indicator.

Yellow (V⁵⁺, VO₂⁺) → blue (V⁴⁺, VO²⁺) → green (V³⁺) → violet (V²⁺). Zinc in hydrochloric acid is used as the reducing agent.

In the vanadium oxidation states practical, ammonium vanadate(V) is added to 1.0 mol dm^-3 hydrochloric acid and dissolves to form a yellow solution. A piece of zinc is then added to reduce the vanadium.

False.

V²⁺ is readily oxidised by oxygen in air. Stopper the test tube once V²⁺ is formed to prevent contact with air, which can cause reoxidation to V³⁺ (green) or beyond — electrode potential data shows O₂ is capable of oxidising V²⁺ to higher oxidation states.

Vanadate(V) contains vanadium in the +5 oxidation state.

The final product V²⁺ contains vanadium in the +2 oxidation state.

Vanadium shows variable oxidation states of +2, +3, +4 and +5. This is a characteristic property of transition metals.

False.

The green appearance can be due to V³⁺ ions, but it may also arise from a mixture of blue V⁴⁺ and yellow V⁵⁺ species. Care is needed to distinguish between the two.

Zinc reduces V⁵⁺ (VO₂⁺, yellow) first to V⁴⁺ (VO²⁺, blue). As reduction continues, V³⁺ (green) then V²⁺ (violet) are formed in sequence, each requiring an additional electron transfer per vanadium ion.

A number assigned to an atom in a compound or ion that represents the degree of oxidation of that atom, indicating how many electrons it has formally gained or lost relative to the neutral element.

A structural isomer is one of two or more compounds that have the same molecular formula but different structural arrangements of atoms.

To determine an unknown concentration by colorimetry, solutions of known concentrations are prepared first and their absorbance measured to plot a calibration curve. The absorbance of the unknown is then read off this curve.

False.

Hexaaquacopper(II) solution absorbs red light (the complementary colour). The blue light is transmitted, which is why the solution appears blue.

1. Prepare metal aqua ion solutions of known concentrations

2. Measure their absorbance and plot a calibration curve (absorbance vs concentration)

3. Measure the absorbance of the unknown solution

4. Read the corresponding concentration from the calibration curve

A colorimeter uses a lamp as a source of white light. Light passes through a filter to produce light of one colour, then through the sample in a cuvette before reaching the detector.

True.

Thiocyanate ions (SCN^-) form intensely coloured complexes with some metal ions (e.g., iron(III) gives a blood-red colour), making very dilute solutions measurable by colorimetry.

The gradient represents the rate of increase in absorbance per unit increase in concentration. A steeper gradient means a greater change in absorbance per unit concentration, indicating a more sensitive measurement for that particular solution and path length.

Autocatalysis is a reaction where one of the products acts as a catalyst for the same reaction, causing the rate to increase as the reaction proceeds before slowing again as reactants are consumed.

In the manganate(VII)/ethanedioate autocatalysis experiment, the product Mn²⁺ catalyses the reaction. The rate is initially slow, then increases, then slows as manganate(VII) is depleted.

True.

Adding the sample to KI solution causes unreacted MnO₄^- to oxidise I^- to I₂, effectively consuming the manganate(VII) and stopping further reaction between MnO₄^- and ethanedioate in that sample.

The MnO₄⁻ in each sample oxidises a fixed amount of I⁻ to I₂, and this I₂ is then titrated with thiosulfate. More MnO₄⁻ means more I₂ produced, requiring more thiosulfate. The titre is directly proportional to [MnO₄⁻].

In the colorimetry method for autocatalysis, manganate(VII) is coloured (purple) and therefore can be monitored using a colorimeter. All other reactants and products are colourless (or too faintly coloured to be detected), so they do not interfere with the absorbance measurement.

True.

Six 10 cm³ samples are removed at 30 s intervals over 180 seconds, giving six time points for the rate analysis.

The rate starts low, rises to a maximum, then falls back to zero. Initially the catalyst (Mn²⁺) concentration is zero so the rate is slow; as more Mn²⁺ is produced the rate increases; the rate then falls as the concentration of MnO₄⁻ is depleted.

Add sodium hydroxide solution and gently heat. Ammonia gas is released, which turns damp red litmus paper blue. The gas has a distinctive pungent smell.

To test for halide ions, nitric acid and then silver nitrate solution are added. A white precipitate indicates Cl^-, a cream precipitate indicates Br^-, and a yellow precipitate indicates I^-.

True.

HCl acidifies the solution and destroys any carbonate ions. Without this step, CO₃^2- would also form a white precipitate with Ba²⁺, giving a false positive result for sulfate.

Carbon dioxide gas is produced and bubbled into limewater (calcium hydroxide solution). The limewater turns milky (cloudy) due to the formation of insoluble calcium carbonate.

Group 2 ions are tested by adding sodium hydroxide solution. Mg(OH)₂ and Ca(OH)₂ form white precipitates. Solubility of Group 2 hydroxides increases down Group 2 from Mg to Ba.

False.

Red litmus paper (not blue) is used to test for hydroxide ions. The presence of OH⁻ makes the solution alkaline, turning red litmus paper blue.

The chemical tests for ions are all carried out in aqueous solution. The ions must be free and mobile in solution to react with the test reagents and produce the characteristic colour changes or precipitates.

A precipitate is an insoluble solid that forms and separates from a solution when two solutions are mixed and a chemical reaction produces a product that exceeds its solubility in the solvent.

Add ammonia solution and observe which precipitate dissolves:

• AgCl (white): dissolves in dilute ammonia

• AgBr (cream): dissolves in concentrated ammonia only

• AgI (yellow): does not dissolve in ammonia

Solubility of silver halides in ammonia decreases from AgCl to AgI as lattice enthalpy increases and the halide becomes more stable.

Sodium hydroxide (NaOH) solution. Adding NaOH causes deprotonation of the coordinated water ligands, forming Cu(OH)₂ (s), a blue precipitate.

In the metal-aqua ion practical, 10 drops of metal-aqua ion solution are added to a test tube, then 10 drops of base. A further 10 drops of base are added to check for changes on addition of excess base.

False.

Some colour changes develop over several minutes. Test tubes should not be disposed of immediately but returned to for a further look, as delayed reactions can provide additional identification information.

An initial white precipitate of Al(OH)₃ forms. On addition of excess NaOH, the precipitate dissolves to give a colourless solution, as Al(OH)₃ is amphoteric and reacts with excess alkali.

In the metal-aqua ion practical, three bases are tested: sodium hydroxide solution, ammonia solution, and sodium carbonate solution.

False.

The iron(III) hydroxide precipitate does not dissolve in excess ammonia. Unlike copper(II) and cobalt(II), Fe³⁺ does not form stable ammonia complexes, so the precipitate remains.

Observations must distinguish between solutions (clear coloured liquids) and precipitates (cloudy solids). Describing both the colour and the physical state is essential for unambiguous identification of the product.

A complex ion formed when a metal cation is surrounded by water molecules acting as ligands, donating lone pairs to the metal centre. For example, [Cu(H₂O)₆]²⁺ is the copper(II) hexaaqua ion.