Edexcel IAL Chemistry A2 Topic 14 Acid–base equilibria

Pearson Edexcel IAL Chemistry syllabusChemistry YCH11First assessment 2019

Write the acid equilibrium first, identify the justified approximation and carry concentration units and logarithms carefully before interpreting acid strength or pH.

Exam points

Write proton-transfer equations and label Brønsted–Lowry conjugate acid–base pairs.
Calculate pH, hydrogen or hydroxide concentration using strong-acid, Ka or Kw relationships.
Compare acid strength through dissociation, Ka and pKa rather than concentration alone.

[Maximum number: 3]

A graph of pH against volume of acid added for an acid-base titration is shown.

(a)

Which acidic solution was used in the titration?

$0.1 \mathrm{~mol} \mathrm{dm}^{-3} \mathrm{CH}_{3} \mathrm{COOH}$

$1.0 \mathrm{~mol} \mathrm{dm}^{-3} \mathrm{CH}_{3} \mathrm{COOH}$

$0.1 \mathrm{~mol} \mathrm{dm}^{-3} \mathrm{HCl}$

$1.0 \mathrm{~mol} \mathrm{dm}^{-3} \mathrm{HCl}$

[ 1 ]

(b)

Which basic solution was used in the titration?

$\mathrm{NH}_{3}$

LiOH

$\mathrm{Ba}(\mathrm{OH})_{2}$

NaOH

[ 1 ]

(c)

A student suggested five indicators that might be used in this titration:
thymol blue
methyl orange
bromophenol blue
bromocresol green
phenolphthalein
How many of these indicators would be suitable? Use your Data Booklet.

[ 1 ]

[Maximum number: 2]

This question is about weak acids.
$\mathrm{p} K_{\mathrm{a}}$ of ethanoic acid, $\mathrm{CH}_{3} \mathrm{COOH}=4.8$ $\mathrm{p} K_{\mathrm{a}}$ of chloroethanoic acid, $\mathrm{CH}_{2} \mathrm{ClCOOH}=2.9$

(a)

What is the pH of a $0.100 \mathrm{~mol} \mathrm{dm}^{-3}$ solution of chloroethanoic acid?

0.27

1.95

2.90

3.90

[ 1 ]

(b)

Which is the acid-conjugate base pair in the reaction between ethanoic acid and chloroethanoic acid?

Acid

Conjugate base

$\mathrm{CH}_{3} \mathrm{COOH}$

$\mathrm{CH}_{3} \mathrm{COO}^{-}$

$\mathrm{CH}_{3} \mathrm{COOH}$

$\mathrm{CH}_{3} \mathrm{COOH}_{2}^{+}$

$\mathrm{CH}_{2} \mathrm{ClCOOH}$

$\mathrm{CH}_{2} \mathrm{ClCOO}^{-}$

$\mathrm{CH}_{2} \mathrm{ClCOOH}$

$\mathrm{CH}_{2} \mathrm{ClCOOH}_{2}^{+}$

[ 1 ]

[Maximum number: 3]

A titration was carried out by adding $0.1 \mathrm{~mol} \mathrm{dm}^{-3}$ hydrochloric acid to $0.1 \mathrm{~mol} \mathrm{dm}^{-3}$ aqueous ammonia.

\mathrm{HCl}(\mathrm{aq})+\mathrm{NH}_{3}(\mathrm{aq}) \rightarrow \mathrm{NH}_{4} \mathrm{Cl}(\mathrm{aq})

The titration curve is shown.

(a)

Which region of the graph represents the most effective buffer solution?

region T

region U

region V

region W

[ 1 ]

(b)

Which of these is the best indicator to use in this titration?
[Refer to the Data Booklet]

methyl red

phenol red

phenolphthalein

thymol blue

[ 1 ]

(c)

What is the approximate pH of an ammonium chloride solution?

2.0

5.8

9.7

11.3

[ 1 ]

[Maximum number: 10]

This question is about sulfuric acid and its salts.

(a)

Sulfur trioxide is used to produce sulfuric acid.

[ 3 ]

(i)

The pH of a $0.10 \mathrm{~mol} \mathrm{dm}^{-3}$ solution of sulfuric acid at $25^{\circ} \mathrm{C}$ is 0.97 .

Calculate the concentration of hydrogen ions, in $\mathrm{mol} \mathrm{dm}^{-3}$ , in this solution.

[ 1 ]

(ii)

In an aqueous solution of sulfuric acid, the following equilibria exist.

\begin{array}{ll} \mathrm{H}_{2} \mathrm{SO}_{4}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \rightleftharpoons \mathrm{H}_{3} \mathrm{O}^{+}(\mathrm{aq})+\mathrm{HSO}_{4}^{-}(\mathrm{aq}) & K_{\mathrm{a}} \text { very large } \\ \mathrm{HSO}_{4}^{-}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \rightleftharpoons \mathrm{H}_{3} \mathrm{O}^{+}(\mathrm{aq})+\mathrm{SO}_{4}^{2-}(\mathrm{aq}) & K_{\mathrm{a}}=0.012 \mathrm{moldm}^{-3} \end{array}

Explain, in terms of these equilibria, why the concentration of hydrogen ions in a $0.10 \mathrm{~mol} \mathrm{dm}^{-3}$ solution of sulfuric acid is not $0.20 \mathrm{~mol} \mathrm{dm}^{-3}$ . No calculation is required.

[ 2 ]

(b)

A buffer solution is made from $\mathrm{HSO}_{4}^{-}$ and $\mathrm{SO}_{4}^{2-}$ ions.

[ 7 ]

(i)

Write two ionic equations involving $\mathrm{HSO}_{4}^{-}$ and $\mathrm{SO}_{4}^{2-}$ ions to show how this

[ 2 ]

(ii)

A buffer solution is formed by mixing
$25.0 \mathrm{~cm}^{3}$ of a solution that is $0.150 \mathrm{~mol} \mathrm{dm}^{-3}$ with respect to $\mathrm{SO}_{4}^{2-}$ ions with $75.0 \mathrm{~cm}^{3}$ of a solution that is $0.100 \mathrm{~mol} \mathrm{dm}^{-3}$ with respect to $\mathrm{HSO}_{4}^{-}$ ions.

Calculate the pH of this buffer solution.
$\left[K_{\mathrm{a}}\right.$ for $\mathrm{HSO}_{4}^{-}$ ions $=0.012 \mathrm{~mol} \mathrm{dm}^{-3}$ ]

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t

[ 5 ]

[Maximum number: 5]

This question is about compounds with the molecular formula $\mathrm{C}_{6} \mathrm{H}_{12} \mathrm{O}_{2}$ .

(a)

Hexanoic acid, $\mathrm{C}_{5} \mathrm{H}_{11} \mathrm{COOH}$ , is a weak acid.

[ 5 ]

(i)

Write the equation for the acid dissociation constant, $K_{a}$ , of hexanoic acid.

[ 1 ]

(ii)

Calculate the pH of a $0.100 \mathrm{~mol} \mathrm{dm}^{-3}$ solution of hexanoic acid.
$\left[\mathrm{p} K_{\mathrm{a}}\right.$ of hexanoic acid $\left.=4.88\right]$

[ 4 ]

[Maximum number: 3]

The compound lactic acid can be synthesised from ethanal in two steps.

(a)

Sodium hydrogencarbonate is part of a buffer in the body that controls the pH of blood. Two of the equilibria involved in this process are shown.

Equilibrium $1 \quad \mathrm{HCO}_{3}^{-}+\mathrm{H}_{3} \mathrm{O}^{+} \rightleftharpoons \mathrm{H}_{2} \mathrm{CO}_{3}+\mathrm{H}_{2} \mathrm{O}$
Equilibrium $2 \quad \mathrm{H}_{2} \mathrm{CO}_{3} \rightleftharpoons \mathrm{CO}_{2}+\mathrm{H}_{2} \mathrm{O}$

[ 3 ]

(i)

Use the equilibria to explain how the buffer keeps the pH of blood nearly constant when a small increase in the concentration of hydrogen ions occurs.
(3)

[ 3 ]

(ii)

The pH of a blood sample was found to be 7.41.

Calculate the ratio of the concentration of $\mathrm{HCO}_{3}^{-}$ to $\mathrm{H}_{2} \mathrm{CO}_{3}$ in the blood sample.

\begin{aligned} & \mathrm{H}_{2} \mathrm{CO}_{3}+\mathrm{H}_{2} \mathrm{O} \rightleftharpoons \mathrm{HCO}_{3}^{-}+\mathrm{H}_{3} \mathrm{O}^{+} \\ & K_{\mathrm{a}}=4.50 \times 10^{-7} \mathrm{~mol} \mathrm{dm}^{-3} \end{aligned}

SECTION C

DO NOTT WRITE IN THIS AREA	DO NOT WRITE IN THIS AREA	DO NOT WRITE IN THIS AREA

Answer ALL the questions. Write your answers in the spaces provided.

[Maximum number: 4]

The equation for the formation of ammonia in the Haber Process is shown

1 / 2 \mathrm{~N}_{2}(\mathrm{~g})+1 \frac{1}{2} \mathrm{H}_{2}(\mathrm{~g}) \rightleftharpoons \mathrm{NH}_{3}(\mathrm{~g})

(a)

Ammonia from the Haber Process reacts with acids.

With phosphoric acid, $\mathrm{H}_{3} \mathrm{PO}_{4}$ , a number of products are formed in solution. One of these is the fertiliser diammonium hydrogenphosphate.

[ 4 ]

(i)

Write an ionic equation to show that ammonium ions are acidic in aqueous solution.
State symbols are not required.

[ 1 ]

(ii)

A solution containing both ammonia and ammonium ions acts as a buffer. Explain, using a relevant ionic equation, the effect of adding a small amount of acid to this buffer.

[ 3 ]

Exam points

Question 4

Question 4(a)

Question 4(b)

Question 4(c)

Question 9

Question 9(a)

Question 9(b)

Question 10

Question 10(a)

Question 10(b)

Question 10(c)

Question 18

Question 18(b)

Question 18(b)(ii)

Question 18(b)(iii)

Question 18(c)

Question 18(c)(i)

Question 18(c)(ii)

Question 20

Question 20(a)

Question 20(a)(i)

Question 20(a)(ii)

Question 19

Question 19(c)

Question 19(c)(i)

Question 19(c)(ii)

Question 22

Question 22(e)

Question 22(e)(ii)

Question 22(e)(iii)