Acid-base & Redox Reactions (OCR A Level Chemistry)

Halide ions can be identified using chemical tests. If an unknown compound is dissolved in dilute nitric acid, and then silver nitrate solution is added, a precipitate will form if the unknown solution contains halide ions. The precipitate formed will be a silver halide.

The general equation for the precipitation reaction of halide ions with silver nitrate solution is:

AgNO₃ (aq) + X^- (aq) → AgX (s) + NO₃^- (aq)

Halide ions can also react with each other in aqueous solutions. Chlorine reacts in a redox reaction with an aqueous solution of sodium bromide, to form sodium chloride and bromine.

Deduce the oxidation state of each of the stated elements in the following compounds and ions

Explain how the oxidation state of the oxygen atom in H₂O₂ is different from its oxidation state in other compounds.

Acid rain can be caused by oxides of nitrogen when they dissolve in water.

Deduce the oxidation state of the nitrogen atom in the following species.

Fossil fuels contain relatively large amounts of sulfur. When these fossil fuels are burnt, sulfur gets into the atmosphere and acts as a pollutant. 

The sulfur is oxidised in the air and rises very high into the atmosphere, where it mixes and reacts with water, oxygen and other chemicals to form more acidic pollutants known as acid rain.

Acid rain causes corrosion of buildings, statues and can damage aquatic and plant life.

Sulfur trioxide reacts with water to form sulfuric acid, H₂SO₄. Write an equation to show the first ionisation of sulfuric acid.

A student reacted sulfuric acid with the weak base ammonia, NH_3. Write an equation for this reaction. 

A student was given homework to plan how to produce a standard solution in a volumetric flask, and then complete a simple titration experiment. 

The student came up with the following method for preparing the standard solution:

• Collect all equipment
• Use a pan balance to measure the amount of solid for the experiment. 
• Tip the solid into a clean volumetric flask and rinse with distilled water to make sure all of the solid is in the flask. 
• Make the solution up to 250 cm³ using distilled water. 
• Add the lid, and shake well to make sure all of the solid is dissolved in the solution.

State four additional instructions which would improve the students’ overall method. These instructions could be placed at any stage of the method. 

The student then completes a titration to find the concentration of sulfuric acid.

They found that 18.50 cm³ of sulfuric acid exactly neutralised 25.00 cm³ of 0.070 mol dm^-3 sodium hydroxide.

Write an equation for this reaction. 

Calculate the concentration of the sulfuric acid, in mol dm^-3.

A solution of 25.0 cm³ ethanoic acid was titrated against 0.150 mol dm^-3 NaOH (aq) and it was found that 22.35 cm³ of the NaOH was needed for complete neutralisation.

Write an equation for the reaction and determine the concentration of the ethanoic acid. 

Ethanoic acid reacts with sodium hydrogen carbonate in an acid base reaction. Write a balanced symbol equation for this reaction. 

Explain why one of the products from this reaction can be described as salt.

A student added calcium to hydrochloric acid and recorded their observations. Write a balanced symbol equation for this reactions and explain using oxidation numbers whether or not this is a redox reaction.

A student wanted to make barium chloride solution from barium carbonate, write an equation to for the formation of barium chloride and suggest one use for this compound.  

Lactic acid is a chemical byproduct of anaerobic respiration, with the chemical formula C₃H₆O₃. Bacteria in yoghurt and in the stomach produce it. Bacteria in beer ferment glucose, producing lactic acid which decreases the pH of the beer and gives it its characteristic sour taste.

A student is required to make up 250 cm³ of an aqueous solution that contains a known mass of lactic acid. The student is provided with a sample bottle containing the lactic acid.

Describe the method, including full apparatus and practical details, that the student should use to prepare this solution. 

Indented content here...The student titrated their standard solution against a solution of sodium hydroxide, NaOH (aq). The burette used by the student for their titration has an error of ± 0.05 cm³. It was used to determine the initial reading, final reading and determining the end point. The student calculated their average titre as 24.50 cm³.

Write an equation for the dissociation of lactic acid.

The concentration of lactic acid is 5.98 x 10^-5 mol dm^-3. The student used 25.0 cm³ samples of this solution to titrate against sodium hydroxide solution.

Calculate the concentration of sodium hydroxide. 

Sulfur dioxide reacts with a solution of copper (II) chloride according to the following unbalanced equation.

SO₂ (g) + _ H₂O (l) + _ CuCl₂ (aq) → H₂SO₄ (aq) + _ HCl (aq) + _ CuCl (s)

Balance the above equation. 

Identify both the specific oxidising agent and reducing agent for the reaction in part a).  

Sodium iodate is one of the main sources of iodine in the world. To extract the iodine, sodium iodate is reacted with sodium hydrogen sulfite, NaHSO₃, according to the following ionic equation.

2IO₃^- (aq) + 5HSO₃^- (aq) → 3HSO₄^- (aq) + 2SO₄^2- (aq) + I₂ (aq) + H₂O (l)

Explain the role of the hydrogen sulfite ions in the reaction shown in part c).

Citric acid, C₆H₈O₇ , is present in lemon juice and is classed as a weak acid.

15.00 cm³ citric acid is reacted with 11.0 g dm^-3 sodium hydroxide, NaOH, to form sodium citrate, Na₃C₆H₅O₇ , and water.

35.80 cm³ of sodium hydroxide was required to react with the lemon juice. 

State the balanced equation for this reaction. 

Calculate the mass, in grams, of sodium hydroxide that reacted with the lemon juice. Give your answer to 3 significant figures. 

Determine the concentration, in mol dm^-3, of citric acid in the sample of lemon juice. Give your answer to 2 significant figures.

During one of the titrations, the student noticed an air bubble in the jet space of the burette.

A student added 2.4 g of a sample of ethanedioic acid, H₂C₂O₄.nH₂O to water and made up the solution to 250 cm³. This solution was titrated against 25.0 cm³ of 0.1 mol dm^-3 KOH solution. The student's solution 

Use the results to calculate the amount, in moles, of ethanedioic acid in the standard solution. 

Using your answer to part a), calculate the following.

At the start of the rough titration, the student had not removed the funnel from the burette. The funnel was removed before the final rough titration reading was taken.

Explain what effect, if any, this will have on the final reading of the rough titration.

The burette used by the student for their titration has an error of ±0.05 cm³. It was used to determine the initial and final reading of the titration.

Calculate the maximum percentage uncertainty in using the burette, using an average titre of 18.70 cm³. Note, this is not the answer for part a). Give your answer to 2 decimal places.

The element magnesium forms a nitrate, Mg(NO₃)₂, which decomposes on heating forming two different gaseous products. 

Write the balanced symbol equation for the decomposition of magnesium nitrate.

Using your answer to part a) and oxidation numbers, explain why the reaction is a redox reaction.

A student heats 4.76 g of Mg(NO₃)₂ and collects the gas at room temperature and pressure, RTP. 

Calculate the total volume of gas, in cm³, produced.

A student carried out a redox titration to find the amount of iron in an iron tablet. They followed the following steps.

1. Crush the iron tablets using the pestle and mortar
2. Transfer the crushed tablets to a weighing boat and measure their combined mass
3. Empty the crushed tablets into the small beaker and reweigh the weighing boat
   Add 100 cm³ of 1.5 mol dm⁻³ of sulfuric acid to the small beaker and stir to dissolve the tablet
4. Filter the solution (to remove any undissolved solids) into the volumetric flask. Rinse the beaker with more sulfuric acid and add the washings to the volumetric flask. Make up to the 250 cm³ mark with distilled water
5. Pipette 25.0 cm³ of this solution into the conical flask
6. Titrate the iron(II) solution with 0.005 mol dm^-3 potassium manganate(VII) solution until the mixture has just turned pink
7. Record your results in an appropriate format and repeat the titration until concordant results are obtained

The unbalanced equation for the reaction is shown below.

__MnO₄^- (aq) + __H⁺ (aq) + __Fe²⁺ (aq) → __Mn²⁺ (aq) + __H₂O (l) + __Fe³⁺ (aq) 

Balance the equation for the titration.

The average titre value was calculated to be 22.35 cm³.

The student used 1.20 g of the iron tablet. 

A commonly accepted percentage of iron in an iron tablet is in the range of 30-35%. Show by calculation, whether the tablets being tested meet this requirement.