Redox equations

Two redox systems are shown below.

Which species is the strongest reducing agent?

A Mg²⁺(aq)

B Mg(s)

C Pb²⁺(aq)

D Pb(s)

This question is about monobasic acids and buffer solutions.

Nitric acid, HNO₃, is a strong monobasic acid.

Calculate the pH of 0.150 mol dm^–3 HNO₃.

Give your answer to 2 decimal places.

Nitrous acid, HNO₂, is a weak monobasic acid.

Predict whether the pH of 0.150 mol dm^–3 HNO₂ is higher, lower or the same as the pH of 0.150 mol dm^–3 nitric acid, HNO₃.

Explain your answer.

Prediction........

Explanation........

When heated, aqueous HNO₂ disproportionates. The equation is shown below.

• What is meant by disproportionation?

• Use oxidation numbers to show that disproportionation has taken place.

* A student plans to prepare a buffer solution with a pH of 4.85 using propanoic acid, CH₃CH₂COOH, and sodium propanoate, CH₃CH₂COONa.

K_a for CH₃CH₂COOH = 1.32 × 10^–5 mol dm^–3.

The student mixes 125 cm³ of 0.180 mol dm^–3 CH₃CH₂COOH(aq) with a calculated mass of CH₃CH₂COONa(s). The student assumes that the volume of the solution does not change.

• Calculate the mass of CH₃CH₂COONa needed to make 125 cm³ of a buffer solution with a pH of 4.85.

• Explain why this solution resists changes in pH when a small amount of alkali is added.

A student sets up a standard cell based on redox systems 2 and 4.

Draw a labelled diagram to show how this cell could be set up to measure its standard cell potential.

Include the conditions required to measure the standard cell potential.

standard conditions........

Calculate the standard cell potential of this cell.

A student predicts that cobalt metal will reduce acidified VO²⁺ ions to form V²⁺ ions in two stages.

Explain why this two-stage reduction should happen in terms of electrode potentials and equilibria shifts.

Include overall equations for the predicted reactions.

This question is about inorganic chemistry.

The flowchart shows reactions involving SO₃^2–(aq) ions.

What is the systematic name of a SO₃^2– ion?

Construct an ionic equation for the reaction of SO₃^2– with H⁺/Cr₂O₇^2–.

The structure of a neutral nickel(II) complex, A, is shown below.

What is the empirical formula of complex A ?

Complete the 3D diagram below to show a stereoisomer of complex A.

This question is about compounds and ions of iron(II) and iron(III) that contain ethanedioate ions, C₂O₄^2–.

The C₂O₄^2– ion, shown below, is an example of a bidentate ligand.

Explain what is meant by the term bidentate ligand.

A complex ion E contains three C₂O₄^2– ions bonded to an iron(III) ion in an octahedral shape.

Complex ion E exists as a mixture of two optical isomers.

Draw 3D diagrams to show the structures of the optical isomers of E.

Include any overall charge.

A student plans an investigation to find the number of waters of crystallisation, x, in a sample of hydrated iron(II) ethanedioate, FeC₂O₄•xH₂O.

The student decides to carry out a redox titration between solutions of iron(II) ethanedioate and potassium manganate(VII) in acidic conditions.

In the titration, both iron(II) ions and ethanedioate, C₂O₄^2–, ions are oxidised.

Construct half-equations for the oxidation of iron(II) and ethanedioate ions.

Oxidation of iron(II) ions

Oxidation of ethanedioate ions

The uncertainty in each burette reading is ±0.05 cm³.

Calculate the percentage uncertainty for the titre in Titration 1.

* In the titration, 5 mol of iron(II) ethanedioate reacts with 3 mol of manganate(VII) ions.

Analyse the student’s results to find the number of waters of crystallisation, x, in the hydrated

iron(II) ethanedioate, FeC₂O₄•xH₂O.

This question is about the chemistry of compounds containing phosphorus.

Phosphorus forms several acids including H₃PO₄ and H₃PO₃.

H₃PO₄ is a tribasic acid. The equilibria for the dissociations are shown below.

During the equilibria, H₂PO₄^– behaves both as an acid and as a base.

Explain this statement, using the equilibria 1, 2 and 3, as required.

In a H₃PO₃ molecule, the O atoms are covalently bonded to the P atom. The H atoms are bonded to the O atoms.

Draw the structure of a H₃PO₃ molecule, showing all the bonds.

On your diagram, add the values for the O–P–O and P–O–H bond angles.

The systematic name of H₃PO₄ is phosphoric(V) acid.

What is the systematic name of H₃PO₃?

Phosphine, PH₃, is a poisonous gas.

Phosphine reacts with oxygen gas to form phosphorus(V) oxide and water.

Write the equation for this reaction.

Aqueous silver nitrate, AgNO₃, is reduced by PH₃. The unbalanced equation is shown below.

Balance the equation and use oxidation numbers to explain why this is a redox reaction.

Explanation........

When phosphorus(V) chloride, PCl₅, and ammonium chloride are heated together, the compound P₃N₃Cl₆ is formed, together with HCl gas.

P₃N₃Cl₆ has a cyclic structure, like the Kekulé structure of benzene.

Write an equation for the reaction of PCl₅ and ammonium chloride to form P₃N₃Cl₆.

Calculate the percentage by mass of P in P₃N₃Cl₆.

Give your answer to 2 decimal places.

Suggest one example of evidence that could show that P₃N₃Cl₆ has a Kekulé structure rather than a delocalised structure.

In a molecule of P₃N₃Cl₆ all the N and Cl atoms are bonded to P atoms.

Suggest a possible structure for a molecule of P₃N₃Cl₆.

This question is about the d-block elements in Period 4 of the periodic table (Sc to Zn).

* Explain, with examples from Period 4, what is meant by the terms d-block element and transition element.

Explain why some d-block elements are not transition elements.

Use electron configurations to support your explanations.

Describe precipitation reactions using either copper or chromium ions as examples.

Include equations.

Describe ligand substitution reactions using either copper or chromium ions as examples.

Include equations.

The ethanedioate ion, C₂O₄^2–, is a bidentate ligand.

A complex ion of cobalt(III) contains two ethanedioate ligands and two water ligands.

Determine the charge of this complex ion and the coordination number of cobalt in the complex ion.

Charge of complex ion........

Coordination number of cobalt........

An acidified solution containing Cr₂O₇^2– ions reacts with vanadium(III) ions in a redox reaction to form a solution containing Cr³⁺ ions and VO₂⁺ ions.

Construct the overall equation for this reaction.

This question is about nitric acid, hydrochloric acid and sulfuric acid.

Dilute nitric acid reacts with aluminium oxide to form a solution of aluminium nitrate.

Write an equation for this reaction.

The solution contains nitrate ions, NO₃^–.

Draw a ‘dot-and-cross’ diagram for the NO₃^– ion.

Use a different symbol for the extra electron.

Balance the half-equation for the oxidation process in this reaction.

In the reduction process in this reaction, HNO₃ and H⁺ react together to form 2 oxides: X (M_r = 30) and Z (M_r = 18).

Determine the formulae of X and Z and write the half-equation for this reduction.

* Three reactions involving sulfuric acid are shown below.

Reaction 1 Dilute sulfuric acid is reacted with nickel(II) hydroxide to form a green solution.

The solvent is allowed to evaporate leaving hydrated crystals of compound D, with the percentage composition by mass: Ni, 22.33%; S, 12.20%; O, 60.87%; H, 4.60%.

Reaction 2 Concentrated sulfuric acid is reacted with hydrogen bromide, HBr, to form three products:

• an element which exists as diatomic molecules

• a gaseous compound E

• a liquid.

At RTP, 1.00 dm³ of compound E has a mass of 2.67 g.

Reaction 3 Concentrated sulfuric acid acts as a catalyst when 2-hydroxypropanoic acid reacts to form compound F (M_r = 144).

In this reaction, 2 mol of 2-hydroxypropanoic acid forms 1 mol of compound F and 2 mol of water.

Identify compounds D, E and F and construct equations for the reactions.

Show structures for any organic compounds.

This question is about reactions of transition metal compounds.

Explain why ethane-1,2-diamine can act as a bidentate ligand.

The iron(III) ion, Fe³⁺, forms a complex ion A with two ethane-1,2-diamine ligands and two chloride ligands.

Complex ion A has cis and trans stereoisomers. One of these stereoisomers exists as optical isomers.

Determine the empirical formula, with charge, of complex ion A and draw the 3-D structures of the three stereoisomers.

Empirical formula with charge........

Structures

Aqueous sodium hydroxide is added to an aqueous solution of iron(II) sulfate. A pale green precipitate forms which turns brown when left to stand in air.

Write an ionic equation for the formation of the pale green precipitate.

* This question is about copper and copper compounds.

Experiment 1 Hydrochloric acid, HCl (aq), is added to an aqueous solution containing [Cu(H₂O)₆]²⁺ complex ions. A yellow-green solution forms containing complex ion B.

Experiment 2 A piece of copper metal is heated with concentrated sulfuric acid. A reaction takes place forming a pale blue solution C and 45 cm³ of a gas D, measured at RTP. The mass of gas D is 0.12 g.

Experiment 3 An excess of copper(II) oxide is heated with dilute nitric acid. The resulting mixture is filtered. The filtrate is a blue solution E. Aqueous potassium iodide, K I (aq), is added to the blue solution E. A white precipitate F and a brown solution G form.

Determine the formulae of B – G. Construct equations for the reactions taking place, include any changes in oxidation number,

and show your working where appropriate.

This question is about redox reactions.

‘Calcium hypochlorite’, Ca(Cl O)₂, is an ionic compound used in ‘bleaching powder’.

The Cl O^– ion in Ca(Cl O)₂ is the active ingredient that kills bacteria.

Calcium hypochlorite is prepared by reacting chlorine gas with calcium hydroxide.

A student analyses the redox reactions shown below. State symbols have been omitted.

The student concludes that different ions containing chromium can act as oxidising or reducing agents.

Using the terms oxidising agent and reducing agent, and ideas about electrode potentials and equilibrium, explain how the student is correct.

Include overall equations.

A student bubbles hydrogen sulfide gas, H₂S(g), through an acidified solution containing manganate(VII) ions, MnO₄^–(aq).

A redox reaction takes place, forming aqueous manganese(II) ions, a yellow precipitate and one other product.

Construct the equation for this reaction. State symbols are not required.

The unbalanced equation for the reaction of silver with concentrated nitric acid is shown below. ….. Ag(s) + ….. NO₃^–(aq) + ….. H⁺(aq) → ….. Ag⁺(aq) +.…. NO(g) +..… H₂O(l) Which numbers for Ag and H₂O will balance the equation?

Your answer

Storage cells and fuels cells are types of electrochemical cell.

The electrode potentials for five redox systems are shown in Table 19.1.

A student sets up an electrochemical cell based on redox systems 1 and 4 in Table 19.1.

Draw a labelled diagram to show how this cell could be set up in the laboratory.

Construct the equation for the overall cell reaction.

In acid conditions, MnO₄^2–(aq) disproportionates to form MnO₂(s) and MnO₄^–(aq).

Explain, in terms of oxidation numbers, why disproportionation has taken place.

Explain, in terms of electrode potentials and equilibrium shifts why MnO₄^2– disproportionates in acid conditions. Use the information in Table 19.1.

An alkaline hydrogen-oxygen fuel cell is set up.

The overall equation for the cell reaction is shown below. 2H₂(g) + O₂(g) → 2H₂O(l) Redox system 2 in Table 19.1 is the positive electrode of this cell.

Write the half-equation at the negative electrode.

The cell potential is 1.23 V.

Calculate the electrode potential of the negative electrode.

State one important feature of a fuel cell that is different from a conventional storage cell.

A student carries out an investigation to identify two metals, M and X, by two different methods.

In Stage 1, a redox reaction takes place between M and HC l (aq), forming hydrogen and a solution containing M in the +2 oxidation state.

Write an overall equation, with state symbols, for this reaction. Write half-equations for the oxidation and reduction processes.

Overall equation........

Oxidation half-equation........

Reduction half-equation........

In Stage 1, suggest two observations that would confirm that all of metal M has reacted.

In Stage 3, write the ionic equation for the reaction taking place in the titration.

Metal M can be identified following the steps below.

1. The amount, in mol, of excess HC l (aq) that remains after the reaction of M with HC l (aq).

2. The amount, in mol, of HC l (aq) that reacted with M.

3. The identity of metal M.

Analyse the results to identify metal M.

The student is provided with the carbonate of an unknown metal, X₂CO₃.

The student measures the mass loss when the X₂CO₃ is reacted with an excess hydrochloric acid. The equation is shown below.

X₂CO₃(s) + 2HCl (aq) → 2XCl (aq) + CO₂(g) + H₂O(l)

The reaction is carried out using this method:

Step 1 Add 100 cm³ HCl (aq) to a conical flask and weigh.

Step 2 Add X₂CO₃ to the conical flask and immediately reweigh.

Step 3 After 5 minutes, reweigh the conical flask and contents.

Calculate the amount, in mol, of CO₂ released in the reaction.

After analysing the results, the student was told that their molar mass of X₂CO₃ was incorrect.

The student evaluated the experiment for possible reasons for the incorrect result.

The student wondered whether the reaction was complete when the mass was recorded after 5 minutes (Step 3).

How could the student modify the experimental procedure to be confident that the reaction was complete?

The student finds out that carbon dioxide is slightly soluble in water.

State and explain how the solubility of CO₂ would affect the calculated molar mass of X₂CO₃.

This question is about reactions of ions and compounds of transition elements.

A student carries out two experiments on a solution containing [Cr(H₂O)₆]³⁺(aq).

Experiment 1 The student adds an excess of aqueous ammonia to a solution containing [Cr(H₂O)₆]³⁺ until a purple solution is formed.

Experiment 2 The student carries out the following reaction sequence.

Step 1 NaOH(aq) is added slowly to a solution containing [Cr(H₂O)₆]³⁺(aq) in a boiling tube. A grey–green precipitate forms.

Step 2 An excess of NaOH(aq) is added to the boiling tube. The precipitate dissolves and a green solution forms containing a 6 coordinate complex ion.

Step 3 H₂O₂ is added to the mixture and the boiling tube is heated. A yellow solution forms.

Step 4 The solution in the boiling tube is acidified. The solution now contains Cr₂O₇²⁻(aq).

What is the formula of the complex ion in the purple solution that forms in Experiment 1

Suggest an equation for the reaction in Experiment 2, Step 1. Include state symbols.

Draw a 3-D diagram for the shape of the complex ion that forms in Experiment 2, Step 2 Include the charge of the ion.

What is the formula of the ion that causes the yellow colour in Experiment 2, Step 3

State the colour of the solution that forms in Experiment 2, Step 4.

Vanadium ions have four common oxidation states. Table 18.1 shows the colours of the ions in aqueous solution.

Complete the electron configuration of a V³⁺ ion.

The student adds excess iron to a solution containing VO²⁺(aq) ions, and observes that the colour of the solution changes from blue to green and then to violet.

Use the relevant standard electrode potentials shown in Table 18.2 to explain these observations.

Construct an equation for the first colour change from blue to green.

Iron(II) gluconate, C₁₂H₂₂FeO₁₄, is the active ingredient in some brands of iron supplements.

A student carries out an experiment to determine the mass of iron(II) gluconate in one tablet of an iron supplement, using the method below.

Stage 1 The student crushes two tablets and dissolves the powdered tablets in dilute sulfuric acid.

Stage 2 The student makes up the solution from Stage 1 to 250.0 cm³ in a volumetric flask.

Stage 3 The student then titrates 25.0 cm³ portions of the solution obtained in Stage 2 with 0.00200 mol dm⁻³ potassium manganate(VII).

The student obtains a mean titre of 13.50 cm³.

In this titration, 1 mol of manganate(VII) ions reacts with 5 mol of iron(II) ions.

Explain why the student used 0.00200 mol dm⁻³ potassium manganate(VII) solution for this titration, rather than the more usual concentration of 0.0200 mol dm⁻³ used in manganate(VII) titrations.

Use the student’s results to determine the mass, in mg, of iron(II) gluconate in one tablet.

Give your answer to 3 significant figures.

This question is about reactions of iron compounds.

In the boxes, write the formulae of A and B.

The student thinks that the reaction of iron(II) sulfide with H⁺ / H₂O is a redox reaction.

Explain, with reasons, whether the student is correct.

Write the equation for the reaction of [Fe(H₂O)₆]²⁺(aq) with Cl₂(g).

Construct an equation for the reaction of H₂S(g) with H⁺(aq)/MnO₄⁻(aq).

This question is about some reactions of d block elements and their ions.

Table 21.1 shows standard electrode potentials which will be needed within this question.

Complete the electron configuration of

a Ni atom: 1s²........

a Ni²⁺ ion: 1s²........

A standard cell is set up in the laboratory with the cell reaction shown below.

Draw a labelled diagram to show how this cell could be set up to measure its standard cell potential.

Include details of apparatus, solutions and the standard conditions required.

Standard conditions........

Predict the standard cell potential of this cell.

Use the information in Table 21.1 to help you answer both parts of this question.

Write the overall equation for the oxidation of Fe²⁺ by acidified H₂O₂.

Zinc reacts with acidified Cr₂O₇^2– ions to form Cr²⁺ ions in two stages.

Explain why this happens in terms of electrode potentials and equilibria.

Include overall equations for the reactions which occur.

* Three different reactions of copper compounds are described below.

Reaction 1: Aqueous copper(II) sulfate reacts with excess aqueous ammonia in a ligand substitution reaction. A deep-blue solution is formed, containing an octahedral complex ion, C, which is a trans isomer.

Reaction 2: Copper(I) oxide reacts with hot dilute sulfuric acid in a disproportionation reaction. A blue solution, D, and a brown solid, E are formed.

Reaction 3: Copper(II) oxide reacts with warm dilute nitric acid in a neutralisation reaction, to form a blue solution. Unreacted copper(II) oxide is filtered off, and the solution is left overnight in an evaporating basin. A hydrated salt, F, crystallises, with the percentage composition by mass: Cu, 26.29%; H, 2.48%; N, 11.59%; O, 59.63%.

Identify C – F by formulae or structures, as appropriate.

Include equations, any changes in oxidation number, and working.

This question looks at reactions of hydrogen peroxide and of cobalt(II) ions.

Aqueous hydrogen peroxide decomposes as shown in equation 2.1.

The reaction is catalysed by manganese(IV) oxide, MnO₂.

A student investigates the decomposition of a hydrogen peroxide solution as outlined below.

• The student adds 50.00 cm³ of H₂O₂(aq) to a conical flask.

• The student adds a small spatula measure of MnO₂ and quickly connects the flask to a gas syringe.

• The student measures the volume of oxygen every 200 seconds.

The student allows the reaction in equation 2.1 to proceed until no more gas is evolved. The volume of O₂ in the syringe is now 55 cm³, measured at RTP.

Calculate the initial concentration of the H₂O₂.

Give your answer to two significant figures.

initial concentration of H₂O₂ =........ mol dm^–3

Hydrogen peroxide can act as an oxidising agent or as a reducing agent.

Some standard electrode potentials are shown below.

Use this information to write an equation for a reaction in which hydrogen peroxide acts as a reducing agent.

Cobalt(II) forms complex ions with water ligands and with chloride ligands.

• With water ligands, cobalt(II) forms a pink octahedral complex ion, [Co(H₂O)₆]²⁺.

• With chloride ligands, cobalt(II) forms a blue tetrahedral complex ion.

A student dissolves cobalt(II) sulfate in water in a boiling tube. A pink solution forms.

Experiment 1 The student places the boiling tube in a water bath at 100 °C. Concentrated hydrochloric acid is added dropwise. The colour of the solution changes from pink to blue.

Experiment 2 The student places the boiling tube from experiment 1 in an ice/water bath at 0 °C. The colour of the solution changes from blue to pink.

Write the equilibrium equation for the reaction that takes place when the colour of the solution changes.

Explain the observations and predict whether the formation of the blue colour is exothermic or endothermic.