Topic 18: Acids and bases

Which species is a Lewis acid but not a Brønsted–Lowry acid?

A.  ${\mathrm{Cu}}^{2+}$

B.  ${{\mathrm{NH}}_4}^{+}$

C.  $\mathrm{Cu}$

D.  ${\mathrm{CH}}_3\mathrm{COOH}$

Which species is a Lewis acid but not a Brønsted–Lowry acid?

A.  ${\mathrm{Cu}}^{2+}$

B.  ${{\mathrm{NH}}_4}^{+}$

C.  $\mathrm{Cu}$

D.  ${\mathrm{CH}}_3\mathrm{COOH}$

In a titration, $25.00 {\mathrm{cm}}^3$ of vinegar required $20.75 {\mathrm{cm}}^3$ of $1.00 \mathrm{mol} {\mathrm{dm}}^{-3}$ potassium hydroxide to reach the end-point.

Calculate the concentration of ethanoic acid in the vinegar.

In a titration, $25.00 {\mathrm{cm}}^3$ of vinegar required $20.75 {\mathrm{cm}}^3$ of $1.00 \mathrm{mol} {\mathrm{dm}}^{-3}$ potassium hydroxide to reach the end-point.

Calculate the concentration of ethanoic acid in the vinegar.

In a titration, $25.00 {\mathrm{cm}}^3$ of vinegar required $20.75 {\mathrm{cm}}^3$ of $1.00 \mathrm{mol} {\mathrm{dm}}^{-3}$ potassium hydroxide to reach the end-point.

Calculate the concentration of ethanoic acid in the vinegar.

Calculate the $K_{\mathrm{b}}$ of the conjugate base of ethanoic acid using sections 2 and 21 of the data booklet.

Calculate the $K_{\mathrm{b}}$ of the conjugate base of ethanoic acid using sections 2 and 21 of the data booklet.

Calculate the $K_{\mathrm{b}}$ of the conjugate base of ethanoic acid using sections 2 and 21 of the data booklet.

Identify the major species, other than water and potassium ions, at these points.

Identify the major species, other than water and potassium ions, at these points.

Identify the major species, other than water and potassium ions, at these points.

Which indicator is appropriate for the acid-base titration shown below?

A. Thymol blue (pK_a = 1.5)
B. Methyl orange (pK_a = 3.7)
C. Bromophenol blue (pK_a = 4.2)
D. Phenolphthalein (pK_a = 9.6)

Which indicator is appropriate for the acid-base titration shown below?

A. Thymol blue (pK_a = 1.5)
B. Methyl orange (pK_a = 3.7)
C. Bromophenol blue (pK_a = 4.2)
D. Phenolphthalein (pK_a = 9.6)

Calculate K_b for HCO₃^– acting as a base.

Calculate K_b for HCO₃^– acting as a base.

Calculate K_b for HCO₃^– acting as a base.

Calculate [H₃O⁺] in the solution and the dissociation constant, K_a , of the acid at 25 °C.

Calculate [H₃O⁺] in the solution and the dissociation constant, K_a , of the acid at 25 °C.

Calculate [H₃O⁺] in the solution and the dissociation constant, K_a , of the acid at 25 °C.

Determine the concentration of methanoic acid in a solution of pH = 4.12. Use section 21 of the data booklet.

Determine the concentration of methanoic acid in a solution of pH = 4.12. Use section 21 of the data booklet.

Determine the concentration of methanoic acid in a solution of pH = 4.12. Use section 21 of the data booklet.

Identify if aqueous solutions of the following salts are acidic, basic, or neutral.

Identify if aqueous solutions of the following salts are acidic, basic, or neutral.

Identify if aqueous solutions of the following salts are acidic, basic, or neutral.

Identify an indicator that could be used for the titration in 5(d)(i), using section 22 of the data booklet.

Identify an indicator that could be used for the titration in 5(d)(i), using section 22 of the data booklet.

Identify an indicator that could be used for the titration in 5(d)(i), using section 22 of the data booklet.

Calculate the pH of 0.100 mol dm⁻³ aqueous ethanoic acid.

K_a = 1.74 × 10⁻⁵

Calculate the pH of 0.100 mol dm⁻³ aqueous ethanoic acid.

K_a = 1.74 × 10⁻⁵

Calculate the pH of 0.100 mol dm⁻³ aqueous ethanoic acid.

K_a = 1.74 × 10⁻⁵

Write an equation to show ammonia, NH₃, acting as a Brønsted–Lowry base and a different equation to show it acting as a Lewis base.

Write an equation to show ammonia, NH₃, acting as a Brønsted–Lowry base and a different equation to show it acting as a Lewis base.

Write an equation to show ammonia, NH₃, acting as a Brønsted–Lowry base and a different equation to show it acting as a Lewis base.

Explain, using appropriate equations, how a suitably concentrated solution formed by the partial neutralization of 2,2-dimethylpropanoic acid with sodium hydroxide acts as a buffer solution.

Explain, using appropriate equations, how a suitably concentrated solution formed by the partial neutralization of 2,2-dimethylpropanoic acid with sodium hydroxide acts as a buffer solution.

Explain, using appropriate equations, how a suitably concentrated solution formed by the partial neutralization of 2,2-dimethylpropanoic acid with sodium hydroxide acts as a buffer solution.

The graph represents the titration of 25.00 cm³ of 0.100 mol dm⁻³ aqueous ethanoic acid with 0.100 mol dm⁻³ aqueous sodium hydroxide.

Deduce the major species, other than water and sodium ions, present at points A and B during the titration.

The graph represents the titration of 25.00 cm³ of 0.100 mol dm⁻³ aqueous ethanoic acid with 0.100 mol dm⁻³ aqueous sodium hydroxide.

Deduce the major species, other than water and sodium ions, present at points A and B during the titration.

The graph represents the titration of 25.00 cm³ of 0.100 mol dm⁻³ aqueous ethanoic acid with 0.100 mol dm⁻³ aqueous sodium hydroxide.

Deduce the major species, other than water and sodium ions, present at points A and B during the titration.

Predict whether the pH of an aqueous solution of ammonium chloride will be greater than, equal to or less than 7 at 298 K.

Predict whether the pH of an aqueous solution of ammonium chloride will be greater than, equal to or less than 7 at 298 K.

Predict whether the pH of an aqueous solution of ammonium chloride will be greater than, equal to or less than 7 at 298 K.

The concentration of excess sodium hydroxide was 0.362 mol dm⁻³. Calculate the pH of the solution at the end of the experiment.

The concentration of excess sodium hydroxide was 0.362 mol dm⁻³. Calculate the pH of the solution at the end of the experiment.

The concentration of excess sodium hydroxide was 0.362 mol dm⁻³. Calculate the pH of the solution at the end of the experiment.

Sketch the neutralisation curve obtained and label the equivalence point.

Sketch the neutralisation curve obtained and label the equivalence point.

Sketch the neutralisation curve obtained and label the equivalence point.

Determine the pH of a 0.250 mol dm⁻³ aqueous solution of ethylamine at 298 K, using section 21 of the data booklet.

Determine the pH of a 0.250 mol dm⁻³ aqueous solution of ethylamine at 298 K, using section 21 of the data booklet.

Determine the pH of a 0.250 mol dm⁻³ aqueous solution of ethylamine at 298 K, using section 21 of the data booklet.

An indicator, HIn, has a pK_a of 5.1.

HIn (aq) $\rightleftharpoons$ H⁺ (aq) + In⁻ (aq)

colour A                      colour B

Which statement is correct?

A.   At pH = 7, colour B would be observed
B.   At pH = 3, colour B would be observed
C.   At pH = 7, [HIn] = [In⁻]
D.   At pH = 3, [HIn] < [In⁻]

An indicator, HIn, has a pK_a of 5.1.

HIn (aq) $\rightleftharpoons$ H⁺ (aq) + In⁻ (aq)

colour A                      colour B

Which statement is correct?

A.   At pH = 7, colour B would be observed
B.   At pH = 3, colour B would be observed
C.   At pH = 7, [HIn] = [In⁻]
D.   At pH = 3, [HIn] < [In⁻]

An aqueous solution containing high concentrations of both NH₃ and NH₄⁺ acts as an acid-base buffer solution as a result of the equilibrium:

NH₃ (aq) + H⁺ (aq) $\rightleftharpoons$ NH₄⁺ (aq)

Referring to this equilibrium, outline why adding a small volume of strong acid would leave the pH of the buffer solution almost unchanged.

An aqueous solution containing high concentrations of both NH₃ and NH₄⁺ acts as an acid-base buffer solution as a result of the equilibrium:

NH₃ (aq) + H⁺ (aq) $\rightleftharpoons$ NH₄⁺ (aq)

Referring to this equilibrium, outline why adding a small volume of strong acid would leave the pH of the buffer solution almost unchanged.

An aqueous solution containing high concentrations of both NH₃ and NH₄⁺ acts as an acid-base buffer solution as a result of the equilibrium:

NH₃ (aq) + H⁺ (aq) $\rightleftharpoons$ NH₄⁺ (aq)

Referring to this equilibrium, outline why adding a small volume of strong acid would leave the pH of the buffer solution almost unchanged.

Suggest a suitable indicator for the titration, using section 22 of the data booklet.

Suggest a suitable indicator for the titration, using section 22 of the data booklet.

Suggest a suitable indicator for the titration, using section 22 of the data booklet.

Sketch a graph of pH against volume of hydrochloric acid added to ammonia solution, showing how you would determine the pK_a of the ammonium ion.

Sketch a graph of pH against volume of hydrochloric acid added to ammonia solution, showing how you would determine the pK_a of the ammonium ion.

Sketch a graph of pH against volume of hydrochloric acid added to ammonia solution, showing how you would determine the pK_a of the ammonium ion.

Explain, using two equations, how an equimolar solution of ammonia and ammonium ions acts as a buffer solution when small amounts of acid or base are added.

Explain, using two equations, how an equimolar solution of ammonia and ammonium ions acts as a buffer solution when small amounts of acid or base are added.

Explain, using two equations, how an equimolar solution of ammonia and ammonium ions acts as a buffer solution when small amounts of acid or base are added.

At 298 K the concentration of aqueous carbon dioxide in carbonated water is 0.200 mol dm⁻³ and the pK_a for Equilibrium (2) is 6.36.

Calculate the pH of carbonated water.

At 298 K the concentration of aqueous carbon dioxide in carbonated water is 0.200 mol dm⁻³ and the pK_a for Equilibrium (2) is 6.36.

Calculate the pH of carbonated water.

At 298 K the concentration of aqueous carbon dioxide in carbonated water is 0.200 mol dm⁻³ and the pK_a for Equilibrium (2) is 6.36.

Calculate the pH of carbonated water.

The reaction of the hydroxide ion with carbon dioxide and with the hydrogencarbonate ion can be represented by Equations 3 and 4.

Equation (3)     OH⁻ (aq) + CO₂ (g) → HCO₃⁻ (aq)
Equation (4)     OH⁻ (aq) + HCO₃⁻ (aq) → H₂O (l) + CO₃²⁻ (aq)

Discuss how these equations show the difference between a Lewis base and a Brønsted–Lowry base.

Equation (3):

Equation (4):

The reaction of the hydroxide ion with carbon dioxide and with the hydrogencarbonate ion can be represented by Equations 3 and 4.

Equation (3)     OH⁻ (aq) + CO₂ (g) → HCO₃⁻ (aq)
Equation (4)     OH⁻ (aq) + HCO₃⁻ (aq) → H₂O (l) + CO₃²⁻ (aq)

Discuss how these equations show the difference between a Lewis base and a Brønsted–Lowry base.

Equation (3):

Equation (4):

The reaction of the hydroxide ion with carbon dioxide and with the hydrogencarbonate ion can be represented by Equations 3 and 4.

Equation (3)     OH⁻ (aq) + CO₂ (g) → HCO₃⁻ (aq)
Equation (4)     OH⁻ (aq) + HCO₃⁻ (aq) → H₂O (l) + CO₃²⁻ (aq)

Discuss how these equations show the difference between a Lewis base and a Brønsted–Lowry base.

Equation (3):

Equation (4):

Aqueous sodium hydrogencarbonate has a pH of approximately 7 at 298 K.

Sketch a graph of pH against volume when 25.0cm³ of 0.100 mol dm⁻³ NaOH (aq) is gradually added to 10.0cm³ of 0.0500 mol dm⁻³ NaHCO₃ (aq).

Aqueous sodium hydrogencarbonate has a pH of approximately 7 at 298 K.

Sketch a graph of pH against volume when 25.0cm³ of 0.100 mol dm⁻³ NaOH (aq) is gradually added to 10.0cm³ of 0.0500 mol dm⁻³ NaHCO₃ (aq).

Aqueous sodium hydrogencarbonate has a pH of approximately 7 at 298 K.

Sketch a graph of pH against volume when 25.0cm³ of 0.100 mol dm⁻³ NaOH (aq) is gradually added to 10.0cm³ of 0.0500 mol dm⁻³ NaHCO₃ (aq).

Which has the strongest conjugate base?

A. HCOOH (K_a = 1.8 × 10⁻⁴)

B. HNO₂ (K_a = 7.2 × 10⁻⁴)

C. HCN (K_a = 6.2 × 10⁻¹⁰)

D. HIO₃ (K_a = 1.7 × 10⁻¹)

Which has the strongest conjugate base?

A. HCOOH (K_a = 1.8 × 10⁻⁴)

B. HNO₂ (K_a = 7.2 × 10⁻⁴)

C. HCN (K_a = 6.2 × 10⁻¹⁰)

D. HIO₃ (K_a = 1.7 × 10⁻¹)

Where is the buffer region for the titration of a weak acid with a strong base?

Where is the buffer region for the titration of a weak acid with a strong base?

Examine the relationship between the Brønsted–Lowry and Lewis definitions of a base, referring to the ligands in the complex ion [CuCl₄]²⁻.

Examine the relationship between the Brønsted–Lowry and Lewis definitions of a base, referring to the ligands in the complex ion [CuCl₄]²⁻.

Examine the relationship between the Brønsted–Lowry and Lewis definitions of a base, referring to the ligands in the complex ion [CuCl₄]²⁻.

Which solutions will form a buffer when mixed?

A.  50 cm³ of 1.0 mol dm⁻³ HCl and 50 cm³ of 1.0 mol dm⁻³ NaOH

B.  50 cm³ of 1.0 mol dm⁻³ CH₃COOH and 50 cm³ of 1.0 mol dm⁻³ NaOH

C.  50 cm³ of 1.0 mol dm⁻³ CH₃COOH and 100 cm³ of 1.0 mol dm⁻³ NaOH

D.  100 cm³ of 1.0 mol dm⁻³ CH₃COOH and 50 cm³ of 1.0 mol dm⁻³ NaOH

Which solutions will form a buffer when mixed?

A.  50 cm³ of 1.0 mol dm⁻³ HCl and 50 cm³ of 1.0 mol dm⁻³ NaOH

B.  50 cm³ of 1.0 mol dm⁻³ CH₃COOH and 50 cm³ of 1.0 mol dm⁻³ NaOH

C.  50 cm³ of 1.0 mol dm⁻³ CH₃COOH and 100 cm³ of 1.0 mol dm⁻³ NaOH

D.  100 cm³ of 1.0 mol dm⁻³ CH₃COOH and 50 cm³ of 1.0 mol dm⁻³ NaOH

A 20.00 cm³ sample of the 0.20 mol dm⁻³ solution of ammonium nitrate is titrated with a 0.20 mol dm⁻³ solution of sodium hydroxide. Determine the pH at the equivalence point, to two decimal places using section 1 and 21 of the data booklet.

A 20.00 cm³ sample of the 0.20 mol dm⁻³ solution of ammonium nitrate is titrated with a 0.20 mol dm⁻³ solution of sodium hydroxide. Determine the pH at the equivalence point, to two decimal places using section 1 and 21 of the data booklet.

A 20.00 cm³ sample of the 0.20 mol dm⁻³ solution of ammonium nitrate is titrated with a 0.20 mol dm⁻³ solution of sodium hydroxide. Determine the pH at the equivalence point, to two decimal places using section 1 and 21 of the data booklet.

Determine the pH of 0.010 mol dm⁻³ 2,2-dimethylpropanoic acid solution.

K_a (2,2-dimethylpropanoic acid) = 9.333 × 10⁻⁶

Determine the pH of 0.010 mol dm⁻³ 2,2-dimethylpropanoic acid solution.

K_a (2,2-dimethylpropanoic acid) = 9.333 × 10⁻⁶

Determine the pH of 0.010 mol dm⁻³ 2,2-dimethylpropanoic acid solution.

K_a (2,2-dimethylpropanoic acid) = 9.333 × 10⁻⁶

Sketch the pH curve for the titration of 25.0 cm³ of ethylamine aqueous solution with 50.0 cm³ of butanoic acid aqueous solution of equal concentration. No calculations are required.

Sketch the pH curve for the titration of 25.0 cm³ of ethylamine aqueous solution with 50.0 cm³ of butanoic acid aqueous solution of equal concentration. No calculations are required.

Sketch the pH curve for the titration of 25.0 cm³ of ethylamine aqueous solution with 50.0 cm³ of butanoic acid aqueous solution of equal concentration. No calculations are required.

The following equation represents the dissociation of water at 25 °C.

2H₂O (l) $\rightleftharpoons$ H₃O⁺ (aq) + OH⁻ (aq)        ΔH = +56 kJ

Which changes occur as the temperature increases?

A. [H₃O⁺] increases and pH will decrease.

B. [H₃O⁺] decreases and pH will increase.

C. [H₃O⁺] increases and pH will increase.

D. [H₃O⁺] decreases and pH will decrease.

The following equation represents the dissociation of water at 25 °C.

2H₂O (l) $\rightleftharpoons$ H₃O⁺ (aq) + OH⁻ (aq)        ΔH = +56 kJ

Which changes occur as the temperature increases?

A. [H₃O⁺] increases and pH will decrease.

B. [H₃O⁺] decreases and pH will increase.

C. [H₃O⁺] increases and pH will increase.

D. [H₃O⁺] decreases and pH will decrease.

Describe, using a suitable equation, how the buffer solution formed during the titration resists pH changes when a small amount of acid is added.

Describe, using a suitable equation, how the buffer solution formed during the titration resists pH changes when a small amount of acid is added.

Describe, using a suitable equation, how the buffer solution formed during the titration resists pH changes when a small amount of acid is added.

Calculate the ratio of [A⁻] : [HA] in a buffer of pH 6.0 given that pK_a for the acid is 4.83, using section 1 of the data booklet.

Calculate the ratio of [A⁻] : [HA] in a buffer of pH 6.0 given that pK_a for the acid is 4.83, using section 1 of the data booklet.

Calculate the ratio of [A⁻] : [HA] in a buffer of pH 6.0 given that pK_a for the acid is 4.83, using section 1 of the data booklet.

Suggest, giving a reason, which point on the curve is considered a buffer region.

Suggest, giving a reason, which point on the curve is considered a buffer region.

Suggest, giving a reason, which point on the curve is considered a buffer region.

Calculate the pH of 0.00100 mol dm^–3 propanoic acid solution. Use section 21 of the data booklet.

Calculate the pH of 0.00100 mol dm^–3 propanoic acid solution. Use section 21 of the data booklet.

Calculate the pH of 0.00100 mol dm^–3 propanoic acid solution. Use section 21 of the data booklet.

Sketch the general shape of the variation of pH when 50 cm³ of 0.001 mol dm^–3 NaOH (aq) is gradually added to 25 cm³ of 0.001 mol dm^–3 CH₃CH₂COOH (aq).

Sketch the general shape of the variation of pH when 50 cm³ of 0.001 mol dm^–3 NaOH (aq) is gradually added to 25 cm³ of 0.001 mol dm^–3 CH₃CH₂COOH (aq).

Sketch the general shape of the variation of pH when 50 cm³ of 0.001 mol dm^–3 NaOH (aq) is gradually added to 25 cm³ of 0.001 mol dm^–3 CH₃CH₂COOH (aq).

Calculate the initial pH before any sodium hydroxide was added, using section 21 of the data booklet.

Calculate the initial pH before any sodium hydroxide was added, using section 21 of the data booklet.

Calculate the initial pH before any sodium hydroxide was added, using section 21 of the data booklet.

Calculate the concentration, in mol dm^–3, of ammonia molecules in the solution with pH = 9.3. Use section 21 of the data booklet.

Calculate the concentration, in mol dm^–3, of ammonia molecules in the solution with pH = 9.3. Use section 21 of the data booklet.

Calculate the concentration, in mol dm^–3, of ammonia molecules in the solution with pH = 9.3. Use section 21 of the data booklet.

Magnesium salts form slightly acidic solutions owing to equilibria such as:

Mg²⁺(aq) + H₂O (l) $\rightleftharpoons$ Mg(OH)⁺(aq) + H⁺(aq)

Comment on the role of Mg²⁺ in forming the Mg(OH)⁺ ion, in acid-base terms.

Magnesium salts form slightly acidic solutions owing to equilibria such as:

Mg²⁺(aq) + H₂O (l) $\rightleftharpoons$ Mg(OH)⁺(aq) + H⁺(aq)

Comment on the role of Mg²⁺ in forming the Mg(OH)⁺ ion, in acid-base terms.

Magnesium salts form slightly acidic solutions owing to equilibria such as:

Mg²⁺(aq) + H₂O (l) $\rightleftharpoons$ Mg(OH)⁺(aq) + H⁺(aq)

Comment on the role of Mg²⁺ in forming the Mg(OH)⁺ ion, in acid-base terms.

Calculate the pH of a 1.00 × 10⁻² mol dm⁻³ aqueous solution of ammonia.

pK_b = 4.75 at 298 K.

Calculate the pH of a 1.00 × 10⁻² mol dm⁻³ aqueous solution of ammonia.

pK_b = 4.75 at 298 K.

Calculate the pH of a 1.00 × 10⁻² mol dm⁻³ aqueous solution of ammonia.

pK_b = 4.75 at 298 K.

Which species can act both as a Lewis acid and a Lewis base?

A.  H₂O

B.  NH₄⁺

C.  Cu²⁺

D.  CH₄

Which species can act both as a Lewis acid and a Lewis base?

A.  H₂O

B.  NH₄⁺

C.  Cu²⁺

D.  CH₄