1.	(a)	(i)	Calculate the concentration, in mol dm–3, of a solution of hydrochloric acid, HCl, which has a pH of 1.13.

(1)
 
(ii)	Calculate the concentration, in mol dm–3 , of a solution of chloric(l) acid, HOCl, which has a pH of 4.23.
Chloric(l) acid is a weak acid with Ka = 3.72 × 10–8 mol dm–3.

(4)
 
(b)	The pH of 0.100 mol dm–3 sulphuric acid is 0.98.
(i)	Calculate the concentration of hydrogen ions, H+, in this solution.

(1)
 
(ii)	Write equations to show the two successive ionisations of sulphuric acid, H2SO4, in water.
......................................................................................................................
......................................................................................................................
(2)

(iii)	Suggest why the concentration of hydrogen ions is not 0.20 mol dm–3 in 0.100 mol dm–3 sulphuric acid.
......................................................................................................................
......................................................................................................................
(1)
 
(c)	Many industrial organic reactions produce hydrogen chloride as an additional product.  This can be oxidised to chlorine by the Deacon process:
4HCl(g) + O2(g)  2Cl2(g) + 2H2O(g)    DH = –115 kJ mol–1.
	0.800 mol of hydrogen chloride was mixed with 0.200 mol of oxygen in a vessel of volume 10.0 dm3 in the presence of a copper(I) chloride catalyst at 400 ºC.  At equilibrium it was found that the mixture contained 0.200 mol of hydrogen chloride.
(i)	Write an expression for the equilibrium constant Kc.

(1)
 
(ii)	Calculate the value of Kc at 400 ºC.

(4)

(d)	State and explain the effect, if any, on the position of equilibrium in (c) of:
(i)	decreasing the temperature;
......................................................................................................................
......................................................................................................................
......................................................................................................................
(2)
 
(ii)	decreasing the volume;
......................................................................................................................
......................................................................................................................
......................................................................................................................
(2)
 
(iii)	removing the catalyst.
......................................................................................................................
......................................................................................................................
......................................................................................................................
(2)
(Total 20 marks)

2.	(a)	What is meant by the term weak acid?
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(2)

(b)	A weak acid, represented by HA, dissociates in water according to the equation:
HA(aq) + H2O(l)  H3O+(aq) + A–(aq)
	Write an expression for the dissociation constant, Ka, for HA.

(1)
 
(c)	25 cm3 of 1.00 mol dm–3 aqueous HA, was titrated with 1.00 mol dm–3 aqueous sodium hydroxide and the pH measured throughout. The titration curve is shown below.

	Use the titration curve to find: 
(i)	the value of the pH at the end point of the titration.
............................................................................................................................
(1)
 
(ii)	the pH of an aqueous solution of the salt NaA. 
............................................................................................................................
(1)
 
(iii)	the value of pKa for the acid HA and, hence the value Ka.
pKa ...................................................................................................................
Ka .....................................................................................................................
(2)
 
(d)	Some of the solutions made during this titration would act as buffer solutions.
(i)	What is meant by the term buffer solution?
............................................................................................................................
............................................................................................................................
............................................................................................................................
(2)
 
(ii)	Use the titration curve to find:
the range of pH values over which this mixture acts as a buffer;
from ....................................................     to ....................................................
(1)
the pH of the most efficient buffer solution.
............................................................................................................................
(1)
 
(e)	Suggest, with reasoning, whether methyl orange or phenolphthalein would be the better indicator for this titration.
Choice ........................................................................................................................
Reasoning ..................................................................................................................
....................................................................................................................................
....................................................................................................................................
(3)

(f)	Explain why, as the titration proceeds, the flask becomes warm but not as warm as it would in a similar titration using 1.00 mol dm–3 solutions of hydrochloric acid and sodium hydroxide.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(3)
 
(g)	A different monobasic weak acid has a dissociation constant of 1.8 × 10– 5 mol dm–3.
(i)	Define pH.
............................................................................................................................
(1)
 
(ii)	Calculate the pH of a 1.00 mol dm–3 aqueous solution of this acid.

 
(3)
(Total 21 marks)

3.	Ethanoic acid, CH3COOH, is a weak acid which can be used, with its salts, to make buffer solutions.
(a)	Explain what is meant by the term weak acid.
....................................................................................................................................
....................................................................................................................................
(1)

(b)	Explain what is meant by the term buffer solution.
....................................................................................................................................
....................................................................................................................................
(2)
 
(c)	An aqueous solution of ethanoic acid of concentration 1.00 mol dm–3 has a pH of 2.8. Sketch, with care, how the pH changes during the titration of 25.0 cm3 1.00 mol dm–3 aqueous ethanoic acid with aqueous sodium hydroxide of the same concentration.

(4)
 
(d)	Indicate on your sketch the portion of the curve where the mixture is behaving as a buffer.
(1)
 
(e)	(i)	Write an expression for the dissociation constant, Ka, for ethanoic acid.

(1)

(ii)	Explain how the pKa of ethanoic could be found from the graph.
............................................................................................................................
............................................................................................................................
(2)
(Total 11 marks)

4.	(a)	The first stage in the manufacture of sulphuric acid is the Contact Process.
2SO2 + O2  2SO3
	At 400 °C the equilibrium constant Kp = 3.00 × 104 atm–1. A catalyst of vanadium(V) oxide is used. In a particular equilibrium mixture at 400 °C the partial pressures of sulphur dioxide and of oxygen were 0.100 atm and 0.500 atm respectively. Show that the yield of SO3 is about 95% of the equilibrium mixture.
(5)
 
(b)	(i)	Pure sulphuric acid is a viscous liquid with a high boiling temperature of
338 °C. It has the structure:

	Suggest in terms of the intermolecular forces in sulphuric acid why it has such a high boiling temperature.
(3)
 
(ii)	Sulphuric acid dissolves in water in a highly exothermic reaction
May her rest be long and placid,
She added water to the acid;
The other girl did what we taught her,
And added acid to the water.
	Suggest why sulphuric acid must always be added to water to dilute it rather than the other way round.
(2)

(c)	Sulphuric acid dissociates in water according to the equations:
H2SO4 + H2O ® H3O+  + HSO
HSO+ H2O   H3O+ + SO
	The dissociation constant for the first dissociation is very large; that for the second is 0.01 mol dm–3 at 25 °C.
 
(i)	Calculate the pH of an aqueous solution containing 0.200 mol dm–3 hydrogen ions.
(1)
 
(ii)	The pH of 0.100 mol dm–3 sulphuric acid is 0.98. Explain why this is so close to the pH of 0.100 mol dm–3 HCl which is 1.0.
(3)
 
(d)	Sulphuric acid is used as the electrolyte in the lead-acid battery found in cars. The electrodes are made from lead and from lead(IV) oxide. As the cell discharges, the lead and the lead(IV) oxide are both converted to lead(II) sulphate, and the sulphuric acid concentration falls.
(i)	Use the information above to deduce the two half equations occurring in the lead acid battery.
(3)
 
(ii)	Hence write an equation to represent the overall process taking place as the cell discharges.
(1)
(Total 18 marks)

5.	(a)	Define the following terms.
(i)	pH .....................................................................................................................
(1)
(ii)	Kw ......................................................................................................................
(1)

(b)	Explain the meaning of the term strong, as applied to an acid or a base.
………………...……………….……………………………………………………..
………………...……………….……………………………………………………..
(1)
 
(c)	Calculate the pH of the following solutions.
(i)	HCl(aq) of concentration 0.200 mol dm–3.

(1)
 
(ii)	NaOH (aq) of concentration 0.800 mol dm–3 (Kw = – 1.00 × 10–14 mol2 dm–6).

(2)
 
(d)	HA is a weak acid with a dissociation constant Ka = 5.62 × 10–5 mol dm–3.
(i)	Write an expression for the dissociation constant, Ka, of HA.

(1)

(ii)	Calculate the pH of a 0.400 mol dm–3 solution of HA.

(3)
 
(e)	A buffer solution contains HA(aq) at a concentration of 0.300 mol dm–3, and its sodium salt, NaA, at a concentration of 0.600 mol dm–3. Calculate the pH of this buffer solution.

(3)
(Total 13 marks)

6.	Propanoic acid is a weak acid which dissociates according to 
CH3CH2COOH(aq)  +  H2O(l)    CH3CH2COO–(aq)  +  H3O+(aq)

 
(a)	(i)	Indicate, in the space provided below the equation, the two acid/base conjugate pairs.
(2)

(ii)	Write the expression for the acid dissociation constant, Ka, for propanoic acid.

(1)
 
(iii)	Calculate the pH of a 0.100 mol dm–3 solution of propanoic acid, for which
Ka = 1.3 × 10–5 mol dm–3.

(3)
 
(iv)	Calculate the concentration of hydroxide ions, OH–, in this same solution of propanoic acid. Kw = 1.00 × 10–14 mol2 dm–6 at the temperature of the solution.

(3)

(b)	If sodium propanoate is dissolved in water, the pH of the resulting solution is not 7, but is near to 8. By writing the equation for the reaction occurring suggest why this is so.
................................................................................................................................
................................................................................................................................
................................................................................................................................
(2)
 
(c)	A mixture of sodium propanoate and propanoic acid behaves as a buffer solution.
(i)	What is meant by a buffer solution?
......................................................................................................................
......................................................................................................................
......................................................................................................................
(2)
 
(ii)	Calculate the pH of a buffer solution that is made by mixing equal volumes of 0.0500 mol dm–3 propanoic acid and 0.100 mol dm–3 sodium propanoate.

(3)
(Total 16 marks)

7.	(a)	The equation below shows the reaction which occurs when ammonia is dissolved in water.
NH3(g) + H2O(1)  NH (aq) + OH–(aq)
(i)	Explain why water is classified as an acid in this reaction.
..........................................................................................................................
..........................................................................................................................
(1)
 
(ii)	The ammonia is acting as a weak base in this reaction.
	What is the difference between a weak base and a strong base?
..........................................................................................................................
..........................................................................................................................
..........................................................................................................................
 (1)
 
(b)	Ammonia reacts with oxygen to form the gases nitrogen(II) oxide and steam.
(i)	Complete the Hess cycle below so that ΔHreaction can be calculated using standard enthalpy changes of formation. Include state symbols.

(2)

(ii)	Calculate ΔHreaction for this reaction using the following data.
ΔHf [NH3(g)] = – 46.1 kJ mol–1
ΔHf [NO(g)] = + 90.2 kJ mol–1
ΔHf [H2O(g)] = – 241.8 kJ mol–1
	Include a sign and units in your answer and give your answer to three significant figures.

(3)
(Total 7 marks)

8.	A saturated solution of calcium hydroxide, Ca(OH)2(aq), has a pH of 9.6.
(a)	Write an expression linking hydrogen ion concentration and pH. Use this to calculate the concentration of hydrogen ions in this solution.

(3)

(b)	(i)	The ionisation constant for water, Kw = 1.0 × 10–14 mol2 dm–6.
	Write the expression for Kw.

	Kw =

(1)
 
(ii)	Calculate the concentration of hydroxide ions in the saturated solution of calcium hydroxide.

(1)
 
(iii)	Calculate the concentration of calcium hydroxide in the saturated solution.

(1)
 
(iv)	Calculate the solubility of calcium hydroxide in g dm–3.
	Give your answer to three significant figures.

(1)

(v)	Suggest why your calculated value may differ significantly from the value in chemistry reference books.
.........................................................................................................................
.........................................................................................................................
(1)
 
(c)	An alternative method for finding the solubility of calcium hydroxide is to titrate 100 cm3 of the saturated solution with hydrochloric acid of concentration 0.00100 mol dm–3.
Ca(OH)2 (aq) + 2HCl (aq) ® CaCl2 (aq) + 2H2O (l)
(i)	Calculate the pH of the hydrochloric acid.

(1)
 
(ii)	Use your answer to (b)(iii) and the information above to calculate the volume of hydrochloric acid needed to neutralise 100 cm3 of the saturated calcium hydroxide solution.

(3)

(iii)	Sketch the titration curve for this reaction.

(2)
 
(iv)	Suggest why phenolphthalein is not a suitable indicator for this reaction.
..........................................................................................................................
..........................................................................................................................
(1)
(Total 15 marks)

9.	(a)	Ammonia reacts with water as below:
NH3(aq)   +   H2O(l)      NH4+(aq)   +   OH–(aq)
	A 0.100 mol dm–3 solution of ammonia has a pH of 11.13.
(i)	Identify the Bronsted–Lowry acid/base conjugate pairs in the equation. Clearly label which are acids and which are bases.
.....................................................................................................................................
.....................................................................................................................................
(2)

(ii)	Draw, on the axes below, a graph to show how the pH of the solution varies as 40 cm3 of 0.100 mol dm–3 hydrochloric acid (a strong acid) is added slowly to 20 cm3 of the ammonia solution.

(4)
 
(iii)	Select, from the following list, the indicator which would be the most suitable for this titration. Give a reason for your choice.
Indicator	pKind	Range	
methyl red	5.1	4.2–6.3	
bromothymol blue	7.0	6.0–7.6	
phenolphthalein	9.3	8.2–10.0	
Indicator: ....................................................................................................................
Reason: ......................................................................................................................
.....................................................................................................................................
(2)

(b)	Nitrous acid, HNO2, is a weak acid with an acid dissociation constant
Ka = 4.70 × 10–4 mol dm–3 at 4 °C.
HNO2(aq)   +   H2O(l)    H3O+(aq) + NO2–(aq)
(i)	Write the expression for Ka.

(1)
 
(ii)	Calculate the pH of a 0.120 mol dm–3 solution of nitrous acid.

(3)

(iii)	Calculate the pH of a buffer solution made by adding 1.38 g of sodium nitrite, NaNO2, to 100 cm3 of the 0.120 mol dm–3 solution of nitrous acid
(Ka = 4.70 × 10–4 mol dm–3).

(4)
 
(iv)	Suggest why a mixture of nitrous acid and sodium nitrite can act as a buffer solution whereas a solution of sodium nitrite on its own does not.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)
(Total 18 marks)

10.	(a)	Benzoic acid is a weak acid. What is meant by a weak acid?
....................................................................................................................................
....................................................................................................................................
(1)
 
(b)	Write an expression for the acid dissociation constant, Ka, for benzoic acid.
(1)

(c)	Ka for benzoic acid has value 6.3 × 10–5 mol dm–3. Use this data to calculate the pH of a solution of benzoic acid of concentration 0.050 mol dm–3.

(3)
 
(d)	In a titration, 40.0 cm3 of 0.050 mol dm–3 sodium hydroxide was added from a burette to 20.0 cm3 of 0.050 mol dm–3 benzoic acid. Sketch a curve on the grid below to show how the pH of the solution would change as the sodium hydroxide was added. Detailed calculations are not required.

(3)
 
(e)	Name a suitable indicator to detect the end-point of this titration.
....................................................................................................................................
(1)

(f)	A mixture of benzoic acid and sodium benzoate can act as a buffer solution.
	In what proportion must solutions of benzoic acid and sodium benzoate, of the same concentration, be mixed to produce a buffer solution of pH 4.5?
	You may use the equation pH = –log Ka –log .
	Ka for benzoic acid has value 6.3 × 10–5 mol dm–3.

(2)
 
(g)	Suggest a salt which would make a buffer solution with pH greater than 7 when mixed with aqueous ammonia.
....................................................................................................................................
(1)
(Total 12 marks)

11.	(a)	(i)	Use an equation to define the term pH.
.....................................................................................................................................
(1)
(ii)	Explain how some solutions can have a negative pH.
.....................................................................................................................................
.....................................................................................................................................
(1)

(b)	The concentration of propanoic acid can be found by titrating a sample with standard sodium hydroxide solution.
(i)	Calculate the pH of 0.100 mol dm–3 propanoic acid at 25 °C; the value of the dissociation constant for the acid, Ka, is 1.30 × 10–5 mol dm–3.

(3)
 
(ii)	Sketch with reasonable accuracy the titration curve that you would expect if 25.0 cm3 of 0.100 mol dm–3 propanoic acid were to be titrated with 0.100 mol dm–3 sodium hydroxide solution.

(iii)	What is the significance of the pH of the mixture when 12.5 cm3 of sodium hydroxide had been added to the propanoic acid?
.....................................................................................................................................
(1)
(Total 10 marks)

12.	Butan-1-ol, a primary alcohol, can be oxidised to form the aldehyde, butanal.
(a)	Give the name or formula of an oxidising agent used in this reaction and of the other reagent required.
	Oxidising Agent
.....................................................................................................................................
	Other Reagent
.....................................................................................................................................
(2)
 
(b)	A possible by-product of this reaction is butanoic acid, CH3CH2CH2CO2H, which is classified as a weak acid.
	Explain what is meant by a weak acid.
Acid ............................................................................................................................
.....................................................................................................................................
Weak ...........................................................................................................................
.....................................................................................................................................
(2)
(Total 4 marks)

13.	(a)	Calculate the pH of a solution of sodium hydroxide of concentration 0.600 mol dm–3, assuming it is a strong base.
	The ionisation constant for water, Kw, is 1.00 × 10–14 mol2 dm–6.

(2)
 
(b)	Calculate the pH of a solution of ethanoic acid of concentration 0.600 mol dm–3.
	The acid dissociation constant, Ka, for ethanoic acid is 1.70 × 10–5 mol dm–3.

(3)
 
(c)	A 100 cm3 sample of sodium hydroxide of concentration 0.600 mol dm–3 is mixed with 200 cm3 of ethanoic acid of concentration 0.600 mol dm–3.
(i)	Write an equation for the reaction which occurs.

(1)

(ii)	Calculate the number of moles of unreacted ethanoic acid remaining in the resulting mixture.

(2)
 
(iii)	Calculate the concentration, in mol dm-3, of the unreacted ethanoic acid in the resulting mixture.

(1)
 
(iv)	The mixture which forms is a buffer. Why does the pH of the mixture remain constant when small quantities of solutions containing H+ or OH– ions are added?
	You may find it helpful to use equations in your explanation.
	Addition of solution containing H+ ions. ...........................................................
.............................................................................................................................
.............................................................................................................................
	Addition of solution containing OH– ions. .........................................................
.............................................................................................................................
.............................................................................................................................
(3)

(v)	The concentration of sodium ethanoate in the mixture is 0.200 mol dm-3. Calculate the pH of the mixture, using your answer to (iii).
	You may use the equation
pH = – log Ka – log 

(2) 
(Total 14 marks)

14.	Pentanoic acid, C4H9COOH, is a weak acid with an acid dissociation constant,
Ka = 1.5 × 10–5 mol dm–3.
(i)	What is meant by the term weak in a weak acid?
.....................................................................................................................................
.....................................................................................................................................
(1)
 
(ii)	Write the expression for the Ka of C4H9COOH.

(1)

(iii)	Calculate the pH of a 0.100 mol dm–3 solution of C4H9COOH.

(3)
 
(iv)	On the grid below sketch the change in pH during the addition of 50.0 cm3 of 0.100 mol dm–3 sodium hydroxide solution to 25 cm3 of 0.100 mol dm–3 pentanoic acid solution.

(4)

(v)	Suggest, with reasoning, a suitable indicator for the titration in (iv).
Indicator	pKind	
Bromophenol blue	4.0	
Methyl red	5.1	
Thymol blue	8.9	
Alizarin yellow	12.5	
Indicator .....................................................................................................................
Reason ........................................................................................................................
.....................................................................................................................................
(2)
(Total 11 marks)

15.	This question is about butanoic acid, CH3CH2CH2CO2H.
(a)	Butanoic acid can be described as a weak acid. Explain what is meant by weak as used in this context.
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
(1)

(b)	The acid dissociation constant, Ka, for butanoic acid is 1.50 ´ 10–5 mol dm–3.
(i)	Write the expression for Ka.

(1)
(ii)	Calculate the pH of a 0.0100 mol dm–3 solution of the acid.

(2)
 
(c)	A buffer solution can be made by mixing aqueous solutions of butanoic acid and sodium butanoate.
	Explain what is meant by a buffer solution.
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
(2)
 
(d)	The following equilibrium reactions are present in the buffer solution.
A	CH3CH2CH2CO2H(l)  CH3CH2CH2CO2–(aq) + H+(aq)
B	H+(aq) +OH–(aq)  H2O(l)
(i)	Which TWO of the substances behave as Brønsted–Lowry bases?
...............................................................................................................................
(1)

(ii)	Use the two equations, A and B, to explain how the buffer solution responds to the addition of a solution containing hydroxide ions.
...............................................................................................................................
...............................................................................................................................
...............................................................................................................................
...............................................................................................................................
(2)
 
(iii)	The pH of a buffer solution is given by the equation
pH = -1g Ka - 1g 
	Using this equation, calculate the pH of a buffer solution made by mixing 100 cm3 of 0.0100 mol dm–3 butanoic acid solution with 300 cm3 of 0.00500 mol dm–3 sodium butanoate solution.
[The Ka for butanoic acid is 1.50 × 10–5 mol dm–3.]
(2)
(Total 11 marks)

16.	(a)	The first step in the esterification of ethanoic acid, CH3COOH, by ethanol in the presence of a small quantity of concentrated sulphuric acid, is the reaction
CH3COOH + H2SO4 → CH3COOH2+ + 

	In the space below the equation, identify the two acid base conjugate pairs.
(2)

(b)	Ethanoic acid, CH3COOH, is a weak acid and dissociates in water according to the equation
CH3COOH(aq) + H2O(l)  H3O+(aq) + CH3COO–(aq)
	Its acid dissociation constant, Ka, is
Ka =  = 1.74 × 10–5 mol dm–3 (at 25 °C)
 
(i)	The concentration of a solution of ethanoic acid can be determined by titrating a 25.0 cm3 sample in a conical flask against a standard solution of sodium hydroxide.
	State whether the pH at the end point is less than 7, 7, or more than 7, and hence name a suitable indicator for this titration.
pH at end point ..................................................................................................
Indicator .............................................................................................................
(2)
 
(ii)	Ethanoic acid is only about 1% ionised in dilute solutions. Its enthalpy of neutralisation is –55 kJ mol–1, whereas the enthalpy of neutralisation of a strong acid, such as hydrochloric acid, is –57 kJ mol–1.
	Explain why there is so little difference between these two values.
............................................................................................................................
............................................................................................................................
............................................................................................................................
............................................................................................................................
............................................................................................................................
............................................................................................................................
(3)

(iii)	Calculate the pH of a 0.140 mol dm–3 solution of ethanoic acid, clearly showing the TWO assumptions that you have made.
Calculation

Assumptions
............................................................................................................................
............................................................................................................................
............................................................................................................................
(4)
 
(iv)	To 50.0 cm3 of the solution in (III) , an equal volume of a 0.200 mol dm–3 solution of potassium ethanoate was added. Calculate the pH of the buffer solution obtained.

(3)
(Total 14 marks)

17.	(a)	(i)	Give the formula of the ion which causes an aqueous solution to be alkaline.
...........................................................
(1)
 
(ii)	Explain why a 0.1 mol dm–3 solution of ammonia has a lower pH than a
0.1 mol dm–3 solution of sodium hydroxide.
................................................................................................................................
................................................................................................................................
................................................................................................................................
(1)
 
(b)	Using the letters P, Q and R arrange the following in order of increasing pH
	de-ionised water	0.1 mol dm–3	0.1 mol dm–3
		ethanoic acid	hydrochloric acid
	P	Q	R
..................................................................
(1)
(Total 3 marks)

18.	Maleic acid has the structure shown below.

	It is a dioic acid and is one of a pair of geometric isomers.
(a)	What is the systematic name for maleic acid?
.......................................................................................................................................
(2)

(b)	Write the equation for the reaction of sodium hydroxide with maleic acid, HO2CCH=CHCO2H, to form disodium maleate (state symbols are not required).

(2)
 
(c)	Maleic acid has two acid dissociation constants with the following values
Ka1 = 1.48 × 10–2 mol dm–3	Ka2 = 8.51 × 10–7 mol dm–3
(i)	Explain why maleic acid has two different Ka values.
................................................................................................................................
................................................................................................................................
................................................................................................................................
(2)
 
(ii)	Complete an equation for the first dissociation of maleic acid.
	HO2CCH=CHCO2H     
(1)
 
(iii)	Maleic acid is a weak acid. Calculate the pH of a solution of maleic acid of concentration 0.0100 mol dm–3 using the value of Ka1 but not Ka2.

(2)

(iv)	Both carboxylic acid groups in maleic acid do dissociate. Suggest how your answer in (c)(iii) might be affected if you consider Ka2 in addition to Ka1?
................................................................................................................................
................................................................................................................................
(1)
(Total 10 marks)

19.	The weak acid methanoic acid, HCOOH, sets up the following equilibrium in water at 298 K:
HCOOH(aq) + H2O(l)  H3O+(aq) + HCOO–(aq)
	The acid dissociation constant, Ka, for methanoic acid at 298 K is 1.78 × 10–4 mol dm–3.
 
(a)	A 0.200 mol dm–3 solution of methanoic acid has a pH of 2.2 at 298 K.
	20.0 cm3 of this solution is titrated with 0.100 mol dm–3 sodium hydroxide solution until excess alkali has been added.
	On the grid below, sketch the titration curve you would expect for this reaction.

(4)

(b)	Equal volumes of 0.500 mol dm–3 methanoic acid and 0.250 mol dm–3 sodium methanoate solution are mixed to make a buffer solution.
(i)	Define the term buffer solution.
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
(2)
 
(ii)	Calculate the pH of this buffer solution.

(3)

(iii)	Explain, with the aid of equations, how this mixture acts as a buffer solution.
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
(4)
(Total 13 marks)

20.	Potassium hydroxide solution, KOH(aq), is used to peel peaches commercially. The peaches are placed in a 7% solution (7.00 g per 100 cm3 of solution) for 3 minutes. The peaches are then washed, cut in half and frozen.
(a)	(i)	Calculate the concentration, in mol dm–3, of the potassium hydroxide solution.

(2)

(ii)	Calculate the pH of this solution.
[Kw = 1.00 × 10–14 mol2 dm–6]

(2)
 
(b)	The concentration of the potassium hydroxide solution must not fall below
1.00 mol dm–3. This is checked at intervals by the titration of 25.0 cm3 of the solution with 0.750 mol dm–3 sulphuric acid.
(i)	Write a balanced equation for the complete neutralisation of sulphuric acid with potassium hydroxide.

(1)
 
(ii)	Calculate the lowest acceptable titration value before more potassium hydroxide must be added to the solution.

(2)

(iii)	Sketch a graph showing how the pH changes when 30 cm3 (an excess) of 0.750 mol dm–3 sulphuric acid is added to 25.0 cm–3 of 1.00 mol dm–3 potassium hydroxide.

(4)
 
(iv)	Suggest a suitable indicator to use for this titration. Use your graph to justify your choice.
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(2)

(c)	Suggest a safety check that should be made before the peaches are frozen. Why is this check necessary?
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(2)
(Total 15 marks)

21.	(a)	(i)	Calculate the pH of 0.050 mol dm–3 hydrochloric acid.

(1)
 
(ii)	Calculate the concentration of hydroxide ions, in mol dm–3, in this solution.
At this temperature, Kw = 1.00 × 10–14 mol2 dm–6.

(1)

(b)	Phosphoric(V) acid, H3PO4, is a weak acid, forming the following equilibrium in water:
H3PO4(aq) + H2O(l)  H2PO4–(aq) + H3O+(aq)
(i)	Write an expression for the acid dissociation constant, Ka, for phosphoric(V) acid.

(1)
 
(ii)	Given that a 0.500 mol dm–3 solution of phosphoric(V) acid has a pH of 1.20, calculate the value of Ka, stating its units.
Assume that there is no further dissociation of the H2PO4– ion.

(4)
 
(c)	The H2PO4– ion formed when phosphoric(V) acid is added to water can dissociate further into HPO42–.
H2PO4–(aq)	+	H2O(l)		HPO42–(aq)	+	H3O+(aq)
..................		..................		..................		..................
 
(i)	In the spaces below the equation, identify the acid base conjugate pairs.
(2)

(ii)	Explain why very little dissociation of the H2PO4– ion occurs in solutions of phosphoric(V) acid.
...........................................................................................................................
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...........................................................................................................................
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(1)
 
(d)	The change in pH when 25.0 cm3 of 0.100 mol dm–3 phosphoric(V) acid is titrated with sodium hydroxide solution of the same concentration can be seen on the graph below.

	From the list below, select a suitable indicator for this titration. Justify your choice.
	pKIn	
bromocresol green	4.7	
bromothymol blue	7.0	
phenolphthalein	9.3	
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(2)
(Total 12 marks)

22.	(a)	Methanoic acid, HCOOH, is a weak acid. Explain what is meant by the terms weak and acid.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)
 
(b)	Write a balanced equation, including state symbols, for the reaction between aqueous solutions of methanoic acid and sodium carbonate.
.....................................................................................................................................
(2)
 
(c)	The following equilibrium is set up when methanoic acid dissociates in water:
	HCOOH(aq) + H2O(1)  HCOO–(aq) + H3O+(aq)

(i)	There are two conjugate acid-base pairs in the above equation.
	Identify them by completing the sentences below:
Formula of one acid is	.......................................................... .
The formula of its conjugate base is	.......................................................... .
(1)
Formula of the other acid is	.......................................................... .
The formula of its conjugate base is	.......................................................... .
(1)
 
(ii)	Write the expression for the acid dissociation constant, Ka, for methanoic acid.

(1)
 
(iii)	Calculate the pH of a 0.100 mol dm–3 solution of methanoic acid at 298 K.
[Ka for methanoic acid is 1.60 × 10–4 mol dm–3 at 298 K]

(3)

(d)	A buffer solution is made up by mixing equal volumes of 0.100 mol dm–3 methanoic acid and 0.400 mol dm–3 sodium methanoate.
(i)	Calculate the pH of the buffer solution obtained.

(3)

(ii)	Explain why the pH of this mixture of methanoic acid and sodium methanoate remains almost constant when a small quantity of an acidic solution, containing hydrogen ions, or a small quantity of an alkaline solution, containing hydroxide ions, is added.
	In your explanation include equations to show the effect of adding hydrogen ions and of adding hydroxide ions.
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(3)
(Total 16 marks)

23.	(a)	What is the formula of the ion found in all acidic solutions?
.....................................................................................................................................
(1)
 
(b)	Which of the following substances would form an acidic solution in water?
	A	CH4	B	NaOH	C	HCO2H
D	NH3	E	HNO3
.....................................................................................................................................
(2)
(Total 3 marks)

24.	10.0 cm3 of a solution of butanoic acid, CH3CH2CH2CO2H, of concentration 0.00660 mol dm–3, was titrated with a solution of aqueous ammonia using a pH probe.
The pH was recorded throughout, and the results were plotted as shown below.

(a)	(i)	Using the pH of butanoic acid from the graph, calculate the initial hydrogen ion concentration.

(2)
 
(ii)	Write the expression for the acid dissociation constant, Ka, for an aqueous solution of butanoic acid.

(1)
 
(iii)	Calculate the value of Ka making the usual assumptions. Give your answer to two significant figures.

(2)
 
(b)	(i)	Write an equation for the reaction between butanoic acid and ammonia. State symbols are not required.

(1)
 
(ii)	Name the two compounds, apart from water, which are present in the mixture between X and Y shown on the graph.
...........................................................................................................................
(2)

(iii)	What type of mixture is present between X and Y? What evidence is there for your answer by reference to the graph?
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(2)
 
(iv)	Explain why it is not possible to carry out this titration using an indicator.
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(1)
 
(v)	Use the graph to estimate the end-point of the titration. Hence calculate the concentration of the ammonia solution.

(2)
(Total 13 marks)

25.	(a)	Sketch the titration curve that you would expect if 25.0 cm3 of 0.100 mol dm–3 hydrochloric acid, HCl, is titrated with 40.0 cm3 of 0.100 mol dm–3 ammonia solution, NH3.

(4)
 
(b)	Using your answer to (a), select a suitable indicator for this titration. Put a tick in the appropriate box in the table below.
Indicator	pKInd	()	
thymol blue	1.7		
bromocresol green	4.7		
phenol red	7.9		
phenolphthalein	9.3		
(1)

(c)	Suggest why there is no suitable indicator for the titration of ethanoic acid with ammonia.
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(2)
(Total 7 marks)

26.	This question concerns the reactions of some compounds of nitrogen.
(a)	The ammonium ion reacts with water and behaves as an acid.
	NH4+(aq) + H2O(l)  NH3(aq) + H3O+(aq)
(i)	Identify the TWO conjugate acid-base pairs in the spaces provided.
acid 1 ................................... base 1 ...................................
acid 2 ................................... base 2 ...................................
(1)
 
(ii)	Write the expression for the acid dissociation constant, Ka, of the ammonium ion.

(1)

(iii)	A solution of ammonium chloride has a pH of 5.00 at 25°C.
Ka for the ammonium ion is 5.62 × 10–10 mol dm–3 at 25°C.
	Calculate the concentration of this solution. State any assumptions you have made.

(4)
 
(iv)	Use the following table and your answer from part (iii) to suggest a suitable indicator for the titration of ammonia solution with hydrochloric acid. Justify your answer.
Indicator	pKIn	
thymol blue	1.7	
methyl red	5.1	
phenolphthalein	9.3 	
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(2)
 
(b)	Hydrogen cyanide is a weak acid in aqueous solution.
	Write an equation to show why aqueous solutions of cyanide ions are alkaline.
.....................................................................................................................................
(1)

(c)	Hydrogen cyanide reacts with propanal as follows:
	CH3CH2CHO + HCN → CH3CH2CH(OH)CN
	Propanal is reacted with a solution of potassium cyanide, KCN, containing a little dilute sulphuric acid.
(i)	What type of reaction is this?
...........................................................................................................................
(1)
 
(ii)	Give the mechanism for the reaction.

(3)
 
(iii)	It is important that the pH is neither too acidic nor too alkaline if a good yield of the product is to be obtained. Explain why this is so.
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(2)

(d)	In an investigation of the kinetics of the nucleophilic substitution reaction between 1-chloropropane and potassium cyanide in aqueous ethanolic solution, the reaction was found to be first order with respect to 1-chloropropane and first order with respect to cyanide ions.
(i)	Give the rate equation for the reaction.
...........................................................................................................................
(1)
 
(ii)	Write a mechanism for the reaction that is consistent with this rate equation.

(3)
(Total 19 marks)

27.	(a)	The table below shows the acid dissociation constants, Ka, of three carboxylic acids.
Acid	Structural formula	Ka/ mol dm–3	
Chloroethanoic	CH2ClCO2H	1.3 × 10–3	
Dichloroethanoic	CHCl2CO2H	5.0 × 10–2	
Trichloroethanoic	CCl3CO2H	2.3 × 10–1	

(i)	Write an expression for the acid dissociation constant, Ka, of chloroethanoic acid.

(1)
 
(ii)	Calculate the pH of a 0.0010 mol dm–3 solution of chloroethanoic acid, making the usual assumptions.

(3)
 
(iii)	Which acid would have the lowest pH at a concentration of 0.0010 mol dm–3?
Use both the data and the structure of the acids to justify your answer. No further calculation is required.
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(2)

(b)	Chloroethanoic acid, CH2ClCO2H, reacts with methanol, CH3OH, in the presence of a sulphuric acid catalyst.
(i)	Draw the displayed formula and give the name of the organic product formed.
	Displayed Formula

Name .................................................................................................................
(3)
 
(ii)	What name is given to the functional group formed in this organic product?
...........................................................................................................................
(1)
 
(iii)	What type of reagent is methanol in this reaction? Explain why it is able to behave in this way and describe how it attacks the chloroethanoic acid.
You may find it helpful to draw a diagram.

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(3)

(iv)	How would you convert the organic product of the reaction between chloroethanoic acid and methanol back into the original compounds?
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(2)
(Total 15 marks)

28.	This question is about propanoic acid, CH3CH2COOH.
(a)	Propanoic acid is a weak acid which dissociates as follows
	CH3CH2COOH(aq) + H2O(l)  CH3CH2COO–(aq) + H3O+(aq)
(i)	In the above equation there are two conjugate acid-base pairs.
	Identify them by completing the sentences below
Formula of one acid is ......................................................................................
The formula of its conjugate base is .................................................................
Formula of the other acid is ..............................................................................
The formula of its conjugate base is .................................................................
(2)
 
(ii)	Propanoic acid is a weak acid. Explain what is meant by the term weak acid.
	Weak
...........................................................................................................................
...........................................................................................................................
	Acid
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(2)

(b)	The acid dissociation constant, Ka, for propanoic acid is 1.30 × 10–5 mol dm–3 at 298 K.
(i)	Write the expression for the acid dissociation constant, Ka, for propanoic acid.

(1)
 
(ii)	A solution of propanoic acid has a pH of 3.44 at a temperature of 298 K.
	Calculate the concentration, in mol dm–3, of the propanoic acid solution. Show clearly two assumptions you have made.
	Calculation:

	Assumptions:
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(5)

(c)	A mixture of sodium propanoate and propanoic acid acts as a buffer solution.
(i)	What is meant by a buffer solution?
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(2)
(ii)	Calculate the pH of a buffer solution made by mixing 100 cm3 of 0.0100 mol dm–3 propanoic acid solution with 300 cm3 of 0.00500 mol dm–3 sodium propanoate solution at 298 K.
	[Ka for propanoic acid is 1.30 × 10–5 mol dm–3 at 298 K]

(3)
(Total 15 marks)

29.	Calculate the pH of the buffer solution formed by mixing 10 cm3 of aqueous benzoic acid of concentration 0.010 mol dm–3 with 40 cm3 of aqueous sodium benzoate of concentration 0.020 mol dm–3.
	For benzoic acid, the acid dissociation constant, Ka, is 6.3 × 10–5 mol dm–3.
	You may find it helpful to use the relationship pH = – log Ka – 

(Total 3 marks)

30.	(a)	The values of the ionic product of water, Kw, at two different temperatures are shown in the table below.
Temperature /°C	Kw
/ mol2 dm–6	
25	1.00 × 10–14	
50	5.48 × 10–14	
 
(i)	Write an equation to represent the ionisation of water.
...........................................................................................................................
(1)
 
(ii)	Write the expression for Kw.
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(1)

(iii)	Define the term pH.
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(1)
 
(iv)	Calculate the pH of pure water at 50 °C.

(2)
 
(v)	Explain why pure water at 50 °C is neutral despite the fact that its pH is not 7.
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(1)

(b)	The pH curve shown below was obtained when a 0.100 mol dm–3 solution of sodium hydroxide was added to 25.0 cm3 of a 0.100 mol dm–3 solution of ethanoic acid.

(i)	What volume of sodium hydroxide solution is required to neutralise half of the ethanoic acid in this reaction?
Volume added = ..............................................................cm3
(1)
 
(ii)	Use the graph to determine the pH when the volume of sodium hydroxide you have stated in part (i) has been added.
pH is .....................................................................................
(1)
 
(iii)	Write an expression for the acid dissociation constant, Ka, of ethanoic acid, CH3COOH.

(1)
 
(iv)	Use your answers to parts (ii) and (iii) to determine the value of Ka for ethanoic acid at the temperature of the titration. Give your answer to two significant figures.

(2)
(c)	Phenolphthalein is a suitable indicator for a titration between ethanoic acid and sodium hydroxide solutions whereas methyl orange is not a suitable indicator.
	Explain why this is so.
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(2)
(Total 13 marks)

31.	The enthalpy change for some reactions cannot be determined directly.
	One such reaction is the thermal decomposition of potassium hydrogencarbonate, which in a closed system at 200 °C is an equilibrium reaction.
	2KHCO3  K2CO3 + H2O + CO2
	However, by determining the enthalpy change for the neutralisation of the two potassium salts with hydrochloric acid, ΔH for the reaction above can be found. The equations for the neutralisation reactions are:
	K2CO3 + 2HCl → 2KCl + H2O + CO2	ΔH1
	KHCO3 + HCl → KCl + H2O + CO2	ΔH2
	ΔH1 and ΔH2 for the neutralisation reactions were determined as follows:
•	30 cm3 of 2 mol dm–3 hydrochloric acid (an excess) was placed in a polystyrene cup, and its temperature measured to the nearest 0.1°C.
•	A weighed quantity of the potassium salt (either the carbonate or the hydrogencarbonate) was added to the acid with rapid stirring, and the temperature measured again when the reaction was complete.
	For the neutralisation using potassium carbonate, the results were as follows:
	Amount of potassium carbonate used = 0.0187 mol
Initial temperature = 23.7 °C
Final temperature = 30.1 °C
 
(a)	State Hess’s Law.
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(1)

(b)	Use the data for the neutralisation of potassium carbonate to calculate the value of ΔH1 to two significant figures. Remember to include a sign and units in your answer.
	[Assume that the heat capacity of the solution is 4.18 J g–1 °C–1, and that it has a mass of 30 g.]

(3)
 
(c)	(i)	Show how the two equations for the neutralisation reactions and their ΔH values can be combined to find a value of ΔH for the thermal decomposition of potassium hydrogencarbonate.
	Calculate this enthalpy change using your value for ΔH1 from part (b), given that ΔH2 = + 29.3 kJ mol–1.

(3)

(ii)	Explain why you would need to include the enthalpy of vaporisation of H2O(l), in order to obtain an accurate value of the enthalpy of decomposition of potassium hydrogencarbonate.
	2KHCO3(s)  K2CO3(s) + H2O(g) + CO2(g)
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(1)
 
(d)	State and explain the effect of a decrease in temperature on the value of the equilibrium constant for the decomposition reaction and hence on the composition of the equilibrium mixture.
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(2)
(Total 10 marks)

32.	Human blood plasma is a buffer solution. It partly owes its buffer properties to carbon dioxide, produced by respiration, dissolving in the blood for transportation to the lungs.
	Carbon dioxide dissolves in water establishing the equilibrium
	CO2(aq) + H2O(l)  H+(aq) + HCO3–(aq)
(a)	(i)	Write the expression for Ka for this equilibrium and give its units.

Units ........................
(2)
 
(ii)	What is the relationship between pKa and Ka?

(1)
 
(b)	Explain what is meant by a buffer solution.
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(2)

(c)	HCO3–(aq) can act as an acid or a base.
	CO2(aq) + H2O(l)  H+(aq) + HCO3–(aq)
	In this equilibrium, decide whether HCO3–(aq) is acting as an acid or as a base. Give the reason for your decision.
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(1)
 
(d)	A sprinter had the pH and total carbonate concentration, [CO2 + HCO3–], of his blood plasma measured immediately before and after a race. The following results were obtained.
	pH	Total carbonate concentration, [CO2 + HCO3–]
/ mol dm–3	
Before race	7.4	2.52 × 10–2	
After race	7.3	1.98 × 10–2	
	The pH of a buffer solution is given by the equation
	
	For this equilibrium pKa is 6.5.

(i)	Use this information, together with the data in the table, to calculate  before the race.

(2)
 
(ii)	The concentration of HCO3– is 0.0224 mol dm–3. Calculate the concentration of CO2 before the race.

(1)
 
(iii)	Use your results from (d)(i) and (ii) to complete the table below which will allow you to compare the results before and after the race.
		[HCO3–]
/ mol dm–3	[CO2]
/ mol dm–3	
Before race		0.0224		
After race	0.158	0.0171	0.00270	
	Two hypotheses have been proposed to explain why vigorous exercise results in an increase in blood plasma acidity (from 7.4 to 7.3).

	Greater muscle activity during a race requires:
	either
Hypothesis I	the combustion of larger quantities of glucose, resulting in an increase in dissolved carbon dioxide and hence an increase in acidity.
	or
Hypothesis II	partial oxidation of glucose to lactic acid and hence an increase in acidity.
	State, giving your reasons, which hypothesis is favoured by the data and your calculations.
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(2)
(Total 11 marks)

33.	This question concerns four solutions, A to D. They were prepared by mixing equal volumes of 0.2 mol dm–3 solutions of two different substances. The substances were
A	HCl(aq) and NaOH(aq)
B	HCl(aq) and NaCl(aq)
C	NH3(aq) and NH4Cl(aq)
D	CH3COOH(aq) and CH3CO2Na(aq)
 
	Select, from A to D, the mixture which would:
(a)	have the lowest concentration of hydrogen ions
A
B
C
D
(1)
 
(b)	act as a buffer of pH about 5
A
B
C
D
(1)
 
(c)	have a chloride ion concentration of 0.2 mol dm–3.
A
B
C
D
(1)
(Total 3 marks)

34.	This question concerns the titration of a solution of sodium hydroxide with a solution of hydrochloric acid. As the titration proceeds the pH of the mixture changes.
(a)	What was the pH when 24.95 cm3 of 1.00 mol dm–3 NaOH(aq) had been added to 25 cm3 of 1.00 mol dm–3 HCl(aq)?
A	3
B	6
C	8
D	11
(1)
 
(b)	What was the pH when 25.05 cm3 of 1.00 mol dm–3 NaOH(aq) had been added to 25 cm3 of 1.00 mol dm–3 HCl(aq)?
A	3
B	6
C	8
D	11
(1)
 
(c)	Which one of the following indicators would be most suitable to use to determine the end point of this titration?
		pH range
A	methyl violet	0–1.6
B	universal indicator	3–11
C	thymolphthalein	8.3–10.6
D	alizarin yellow R	10.1–13.0
(1)
(Total 3 marks)

35.	This question is about the reaction of magnesium with hydrochloric acid which takes place rapidly at room temperature.
	Mg(s) + 2HCl(aq) ® MgCl2(aq) + H2(g)	DH = –467 kJ mol–1
 
(a)	Rewrite the equation omitting spectator ions.
(1)
 
(b)	Suggest the sign of the following entropy changes for this reaction. Justify each of your answers.
(i)	DSsystem
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(2)
(ii)	DSsurroundings
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(2)
(iii)	DStotal
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(1)

(c)	A student carried out this experiment at five different temperatures in order to calculate the activation energy of the reaction. The student’s laboratory record is shown below.
Method
Clean a strip of magnesium weighing 0.100 g with sand paper. Measure the temperature of 20 cm3 of 1.00 mol dm–3 hydrochloric acid in a 100 cm3 beaker. Add the magnesium ribbon, stir continuously, and time how long it takes for the magnesium to disappear. Repeat the experiment at four other temperatures.
Assumption: the initial rate of reaction is proportional to 1/time.
Results	
Temperature
/°C	Temperature
/K	1/T
/K–1	time
/s	1/time
/s–1	ln 1/time	
24	297	3.37 × 10–3	45	0.0222	–3.81	
33	306	3.27 × 10–3	25	0.0400	–3.22	
45	318	3.14 × 10–3	11	0.0909	-2.40	
56	329	3.04 × 10–3	6	0.1667	-1.79	
10	283	3.53 × 10–3	122	0.0082	-4.80	
 
	The Arrhenius equation is ln k = –Ea/R × (1/T) + constant
	ln 1/time is proportional to ln k and so a graph of ln 1/time will have the same gradient as that of the Arrhenius plot of ln k against 1/Temperature
	The student plotted the graph of ln 1/time against 1/Temperature and from this the activation energy, EA, was calculated as + 51.3 kJ mol–1.
(i)	Suggest the reason for cleaning the magnesium ribbon with sand paper.
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(1)

(ii)	Calculate the number of moles of hydrochloric acid used up when all the magnesium reacts in one experiment. Hence comment on whether the change in concentration during the reaction will have a significant effect on the validity of the assumption that the initial rate of reaction is proportional to 1/time. How would you overcome this potential error?
	[Take the relative atomic mass of magnesium as 24 in this and subsequent calculations.]
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(5)

(iii)	Use the value of DH and other information given in the question to calculate the temperature change in an experiment assuming no energy is lost to the surroundings. Hence comment on whether this change in temperature will have a significant effect. How would you overcome this potential error?
	[DH = –467 kJ mol–1.
	heat produced = mass × specific heat capacity × change in temperature.
	Assume that the specific heat capacity of the solution is 4.18 J K–1 g–1]

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(4)

(iv)	The most difficult thing to measure accurately is the time it takes for the magnesium to disappear and the time measured can be up to 2 seconds out.
Assuming this error, calculate the shortest time at 56 °C and the longest time at 10 °C for this reaction.
	Complete the table for these times. Plot the two points on the grid below and join them with a straight line. From the gradient, which equals –EA/R, of this line calculate another value for the activation energy.
Temperature
/ °C	Temperature
/K	1/T
/K–1	time
/s 	1/time
/s–1	ln 1/time	
56	329	3.04 × 10–3				
10	283	3.53 × 10–3				

(v)	If the reaction mixture is not stirred, the magnesium tends to float on the surface of the acid.
	Suggest how this would affect the measurements of the rate of the reaction.
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(1)
 
(vi)	Suggest two other improvements the student could do to this experiment to improve the accuracy or validity of the results.
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(2)
 
(vii)	If ethanoic acid of the same concentration and at the same temperature is used instead of hydrochloric acid, explain how the rate would differ.
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(1)
(Total 24 marks)