Iodine reacts with nitric acid to form nitrogen dioxide, NO₂.

Balance the following equation for this reaction.

I₂ + 10HNO₃   ➔   ......HIO₃  +  ......NO₂  +  ......H₂O 

Give the oxidation state of iodine in I₂ and HIO₃.

Two half equations involving iodine are shown below.  Use these two half equations to write an overall equation for the reaction.

IO₃⁻ + 5e⁻ + 6H⁺ ➔ 3H₂O +  $\frac{1}{2}$ I₂ 

I⁻ ➔ $\frac{1}{2}$ I₂  + e⁻ 

Iodide ions can behave as a reducing agent in the reaction with chlorine.  State the meaning of the term 'reducing agent'.

Write the simplest ionic equation for the reaction of chlorine and iodide ions. 

Copper can be extracted from CuS by reaction with nitric acid to produce a solution of copper sulfate, as shown in the reaction below.

3CuS_(s) + 8HNO_3(aq) ➔ 3CuSO_4(aq) + 8NO_(g) + 4H₂O_(l)

Give the oxidation states of nitrogen in HNO₃ and NO.

Deduce a half equation for the reduction of the nitrate ion in acidic conditions to form NO and H₂O.

Deduce a half equation for the oxidation of the sulfide ion to form the sulfate ion and H⁺ ions in solution.

NO can also be formed when sodium nitrite (NaNO₂) is reacted with iodide ions. 

Give the oxidation state of nitrogen in NO₂⁻ and NO, and write a half equation for the conversion of NO₂⁻ to NO.

During this reaction, iodide ions are converted to iodine. 

Write an overall equation for the reaction, and state the role of NO₂⁻ in this reaction.

This question is about the redox reactions of chlorine.

Chlorine and water react in a redox reaction according to the equation below.

Cl₂ + H₂O ⇌ HClO + HCl

Write two half equations to show the redox processes happening in this reaction.

Hydrochloric acid can be oxidised with sodium chlorate (I) (NaClO) solution. Write the following equations:

• A half equation for the oxidation of chloride ions to chlorine
• A half equation for the reduction of chlorate ions (ClO⁻) to chlorine in acidic conditions 
• An overall equation for the reaction 

ClO⁻ ions are found in bleach and their concentration can be found by the reaction with iodide ions.

ClO⁻ + 2I⁻ + 2H⁺ ➔ I₂ + Cl⁻ + H₂O 

10 cm³ of bleach, containing ClO⁻ ions, was reacted with 69.5 mg of potassium iodide. 

Find the concentration of ClO⁻ ions in the bleach solution, giving your answer to 3 significant figures.

The dichromate ion, Cr₂O₇²⁻, can be used to oxidise iodide ions to iodine and sulfite ions to sulfate ions.

What is the oxidation state of chromium in Cr₂O₇²⁻?

Write the following equations:

• A half equation for the oxidation of iodide ions to iodine
• A half equation for the reduction of Cr₂O₇²⁻ to Cr³⁺ in acidic conditions 
• An overall redox equation for the reaction 

Write the following equations:

• A half equation for the oxidation of SO₃²⁻ to SO₄²⁻
• A half equation for the reduction of Cr₂O₇²⁻ to Cr³⁺ in acidic conditions
• An overall equation for the reaction 

Ethanedioic acid, (COOH)₂, is the active ingredient in a spray to treat infections of the outer ear.

A chemist carries out a redox titration using aqueous cerium(IV) sulfate, Ce(SO₄)_2(aq), to determine the percentage, by mass, of ethanedioic acid in the spray.

In the titration, Ce⁴⁺_(aq) ions oxidise ethanedioic acid in acidic conditions, as shown by the equation below

2Ce⁴⁺_(aq) + (COOH)_2(aq) → 2Ce³⁺_(aq) + 2CO_2(g) + 2H⁺_(aq)

Ce⁴⁺_(aq) ions have a yellow colour. Ce³⁺_(aq) ions are colourless.

Dilute sulfuric acid is added to 5.00 g of the spray to form a colourless solution which is made up to 250.0 cm³ with distilled water.

25.00 cm³ of this solution is titrated with 1.00 x 10^-2 moldm^-3 Ce(SO₄)₂ from the burette.

The titration is repeated until concordant titres are obtained. 

The mean titre is 22.25 cm³.

What colour change would the chemist observe at the end point?

Calculate the percentage, by mass, of ethanedioic acid, (COOH)₂, in the spray.

Give your answer to 3 significant figures.

A chemist uses the following procedure to analyse iron tablets containing iron in the form of iron(II) sulfate, FeSO₄.

1. Crush 1 iron tablet and transfer the resulting powder to a 250.0 cm³ volumetric flask. 
2. Rinse the apparatus used with dilute sulfuric acid transferring all washings to the volumetric flask.
3. Add dilute sulfuric acid to make up the solution to exactly 250.0 cm³. Stopper the flask and invert it several times.
4. Titrate 25.00 cm³ of this solution with 2.00 x 10^-3 mol dm^-3 potassium manganate(VII) solution.
5. Repeat the titration until concordant results are obtained.

The overall equation for the reaction occurring in the titration is:

MnO₄^-_(aq) + 8H⁺_(aq) + 5Fe²⁺_(aq) ⇌ Mn²⁺_(aq) + 4H₂O_(l) + 5Fe³⁺_(aq)

The table below shows some standard electrode potential data.

Using data in the table, explain why the chemist used sulfuric acid and not hydrochloric acid in their titration.

Suggest why this titration does not require the addition of an indicator.

State the colour change at the end-point of this titration.

The chemist decides to take burette readings from the top of the meniscus rather than from the bottom. 

Explain the effect of this, if any, on the titre values.

The chemist's titration readings are shown in the table below.

Complete the table.

Use the titration results in part e) and information from the procedure to calculate the mass of FeSO₄ in 1 iron tablet.

Give your answer to 3 significant figures.

Potassium sulfite, K₂SO₃, is used as a preservative in some foods and drinks.

Food safety laws permit a maximum of 0.500 g of K₂SO₃ per kg of sausage meat.

A scientist uses a manganate(VII) titration to determine how much K₂SO₃ a sample of sausage meat contains.

Step 1: the K₂SO₃ from 1.00 kg of sausage meat is extracted to form a solution containing aqueous SO₃^2- ions.

Step 2: the solution from step 1 is made up to 250.0 cm³ in a volumetric flask with water. 25.0 cm³ of this diluted solution is pipetted into a conical flask.

Step 3: the pipetted solution from step 2 is acidified with dilute sulfuric acid and then titrated with 1.20 x 10^-2 mol dm^-3  solution of potassium manganate(VII), KMnO₄.

The equation for this titration is:

2MnO₄^-_(aq) + 6H⁺_(aq) + 5SO₃^2-_(aq) ➔ 2Mn²⁺_(aq) + 3H₂O_(l) + 5SO₄^2-_(aq)

10.40 cm³ of KMnO_4(aq) is required to reach the endpoint.

Analyse the results to determine whether the amount of potassium sulfite in sauage meat is above or below the legal limit.

Compound Y is an iodate(V) salt of a group 1 metal.

The iodate(V) ion has the formula IO₃^-.

A scientist carries out a titration to find the formula of compound Y.

Step 1: the scientist dissolves 1.39 g of Y in water and makes up the solution to 250.0 cm³ in a volumetric flask.

Step 2: the scientist pipettes 20.00 cm³ of the solution of Y into a conical flask, followed by 10 cm³ of dilute sulfuric acid and an excess of KI_(aq).

The iodate(V) ions are reduced to iodine, as shown below.

IO₃^-_(aq) + 6H⁺_(aq) + 5I^-_(aq) ➔ 3I_2(aq) + 3H₂O_(l)

Step 3: the resulting mixture is titrated with 0.200 mol dm^-3 Na₂S₂O_3(aq).

2S₂O₃^2-_(aq) + I_2(aq) ➔ S₄O₆^2-_(aq) + 2I^-_(aq)

The scientist's titration readings are shown in the table below.

Complete the table.

Calculate the mean titre, in cm³, that the scientist should use for analysing the results.

Describe and explain how the student should determine the end point of this titration accurately.

Determine the relative formula mass and formula of the group 1 iodate(V), Y.