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CHEMISTRY, M6 2024 HSC 3 MC

Which of the following compounds can be correctly described as an Arrhenius base when dissolved in water?

  1. Sodium nitrate
  2. Sodium sulfate
  3. Sodium chloride
  4. Sodium hydroxide
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\(D\)

Show Worked Solution
  • An Arrhenius base is a compound that increases the concentration of \(\ce{OH-}\) ions when it is dissolved in water.
  • Sodium hydroxide is the only compound that dissolves in water to produce hydroxide ions.
  •    \(\ce{NaOH(s) -> Na+(aq) + OH-(aq)}\)

\(\Rightarrow D\)

CHEMISTRY, M6 2023 HSC 22

Explain how the following substances would be classified under the Arrhenius and Brønsted-Lowry definitions of acids. Support your answer with relevant equations.  (4 marks)

    •  \( \ce{HCl(aq)} \)
    • \( \ce{NH4Cl(aq)} \)
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  • Acids are defined by Arrhenius as hydrogen-containing compounds that dissociate in water to give \(\ce{H+}\) ions.
  • \(\ce{HCl(aq)}\) produces \(\ce{H+}\) ions in water and therefore qualifies within Arrhenius’ definition of an acid.
  •    \(\ce{HCl(aq) \rightarrow H+(aq) + Cl−(aq)}\)
  • The salt \(\ce{NH4Cl}\) would not be recognised as an acid with Arrhenius’ definition, since the predominant ions present in aqueous solution are ammonium and chloride.
  • The Brønsted−Lowry theory states that acids are proton donors. \(\ce{HCl(aq)}\) is a proton donor and therefore also qualifies as a Brønsted−Lowry acid.
  • In contradiction to Arrhenius, ammonium chloride \(\ce{(NH4Cl)}\) is also classified as a Brønsted−Lowry acid. This is due to the ammonium ion donating a proton to water to form a hydronium ion.
  •    \(\ce{NH4+(aq) + H2O(l) \rightleftharpoons NH3(aq) + H3O+(aq)}\)
Show Worked Solution
  • Acids are defined by Arrhenius as hydrogen-containing compounds that dissociate in water to give \(\ce{H+}\) ions.
  • \(\ce{HCl(aq)}\) produces \(\ce{H+}\) ions in water and therefore qualifies within Arrhenius’ definition of an acid.
  •    \(\ce{HCl(aq) \rightarrow H+(aq) + Cl−(aq)}\)
  • The salt \(\ce{NH4Cl}\) would not be recognised as an acid with Arrhenius’ definition, since the predominant ions present in aqueous solution are ammonium and chloride.
  • The Brønsted−Lowry theory states that acids are proton donors. \(\ce{HCl(aq)}\) is a proton donor and therefore also qualifies as a Brønsted−Lowry acid.
  • In contradiction to Arrhenius, ammonium chloride \(\ce{(NH4Cl)}\) is also classified as a Brønsted−Lowry acid. This is due to the ammonium ion donating a proton to water to form a hydronium ion.
  •    \(\ce{NH4+(aq) + H2O(l) \rightleftharpoons NH3(aq) + H3O+(aq)}\)

CHEMISTRY, M6 EQ-Bank 8 MC

Which of the following is NOT a Bronsted-Lowry reaction?

  1. \(\ce{NH4^+ + NH2^- \rightleftharpoons 2NH3}\)
  2. \(\ce{CO2 + OH^- \rightleftharpoons HCO3^-}\)
  3. \(\ce{HClO4 + CH3COOH \rightleftharpoons CH3COOH2^+ + ClO4^-}\)
  4. \(\ce{CH3CH2O^- + CH3NH3^+ \rightleftharpoons CH3CH2OH + CH3NH2}\)
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`B`

Show Worked Solution
  • A Bronsted-Lowry reaction occurs when one species (acid) transfers a proton to another species (base).
  • Although it is an acid-base reaction, no proton transfer occurs in  \(\ce{CO2 + OH^- \rightleftharpoons HCO3^-}\)

`=> B`

CHEMISTRY, M6 2015 HSC 28

The equipment shown is set up. After some time a ring of white powder is seen to form on the inside of the glass tube.

  1. Why would this NOT be an acid-base reaction according to Arrhenius?   (1 mark)

  2. Explain why this would be considered a Bronsted-Lowry acid-base reaction. Include an equation in your answer.   (2 marks)

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a.   According to Arrhenius:

  • An acid is a solution that produces hydrogen ions when in a solution.
  • A base is a solution that produces hydroxide ions when in a solution.
  • This reaction does not occur in an aqueous solution and would not be an acid-base reaction according to Arrhenius.
     

b.   \(\ce{HCl(g) + NH3(g) -> NH4+ + Cl-}\)

  • A Bronsted-Lowry acid donates a proton while a base accepts a proton.
  • This reaction involves proton transfer (\(\ce{HCl}\) donates, \(\ce{NH3}\) receives) and would therefore be considered a Bronsted-Lowry acid-base reaction.
Show Worked Solution

a.   According to Arrhenius:

  • An acid is a solution that produces hydrogen ions when in a solution.
  • A base is a solution that produces hydroxide ions when in a solution.
  • This reaction does not occur in an aqueous solution and would not be an acid-base reaction according to Arrhenius.

♦♦ Mean mark (a) 37%.

b.   \(\ce{HCl(g) + NH3(g) -> NH4+ + Cl-}\)

  • A Bronsted-Lowry acid donates a proton while a base accepts a proton.
  • This reaction involves proton transfer (\(\ce{HCl}\) donates, \(\ce{NH3}\) receives) and would therefore be considered a Bronsted-Lowry acid-base reaction.

Mean mark (b) 53%.

CHEMISTRY, M6 2019 HSC 28

Assess the usefulness of the Brønsted-Lowry model in classifying acids and bases. Support your answer with at least TWO chemical equations.   (5 marks)

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  • The Bronsted-Lowry model is a way of classifying acids and bases based on their ability to donate or accept protons.
  • This model is more comprehensive than the Arrhenius model, as it can explain the acid-base behaviour of more species, including those that do not contain \(\ce{OH-}\) ions, and non-aqueous acid-base reactions.
  • Consider the reaction \(\ce{NH3(g) + HCl(g) -> NH4Cl(s)}\) where a proton is transferred from hydrogen chloride to ammonia (according to the Bronsted-Lowry model). Ammonia is not an Arrhenius base as it doesn’t dissociate to produce \(\ce{OH-}\) ions and the reaction cannot be described using the Arrhenius model.
  • However, the Bronsted-Lowry model does have some limitation, such as its inability to explain the acidity of certain acidic oxides and their reactions with basic oxides.
  • e.g. \(\ce{CaO(s) + SO3(g) -> CaSO4(s)}\) is an acid-base reaction but since there is no proton transfer, it cannot be described using the Bonsted-Lowry model.
Show Worked Solution
  • The Bronsted-Lowry model is a way of classifying acids and bases based on their ability to donate or accept protons.
  • This model is more comprehensive than the Arrhenius model, as it can explain the acid-base behaviour of more species, including those that do not contain \(\ce{OH-}\) ions, and non-aqueous acid-base reactions.
  • Consider the reaction \(\ce{NH3(g) + HCl(g) -> NH4Cl(s)}\) where a proton is transferred from hydrogen chloride to ammonia (according to the Bronsted-Lowry model). Ammonia is not an Arrhenius base as it doesn’t dissociate to produce \(\ce{OH-}\) ions and the reaction cannot be described using the Arrhenius model.
  • However, the Bronsted-Lowry model does have some limitation, such as its inability to explain the acidity of certain acidic oxides and their reactions with basic oxides.
  • e.g. \(\ce{CaO(s) + SO3(g) -> CaSO4(s)}\) is an acid-base reaction but since there is no proton transfer, it cannot be described using the Bonsted-Lowry model.

♦♦♦ Mean mark 35%.
MARKER’S COMMENT: Two relevant equations required to achieve full marks.

CHEMISTRY, M6 2020 HSC 34

The effect of concentration on the pH of acrylic acid `(text{C}_2 text{H}_3 text{COOH})` and hydrochloric acid `(text{HCl})` solutions is shown in the graph. Both of these acids are monoprotic.
 

 

Explain the trends in the graph. Include relevant chemical equations in your answer.   (4 marks)

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\(\ce{HCl(aq) -> H+ (aq) + Cl– (aq)}\)

\(\ce{C2H3COOH(aq) \rightleftharpoons H+ (aq) + C2H3COO– (aq)}\)

  • \(\ce{HCl}\) is a strong acid that fully dissociates in water, resulting in a high concentration of \(\ce{H+}\) ions and a low pH.
  • Acrylic acid, on the other hand, is a weak acid that only partially dissociates in water, resulting in a lower concentration of \(\ce{H+}\) ions and a higher pH.
  • When the concentration of \(\ce{HCl}\) decreases by a factor of 10, its pH increases by 1 due to the decrease in \(\ce{H+}\).
  • However, when the concentration of acrylic acid decreases by a factor of 10, its pH increases by less than 1.
  • This is due to the change in pH causing the equilibrium to shift right, producing more \(\ce{H+}\) ions in response to the dilution, resulting in a smaller change in the concentration of \(\ce{H+}\), and thus smaller change in pH.
  • At very dilute concentrations, the degree of dissociation of acrylic acid approaches 100% and its pH converges closely to that of \(\ce{HCl}\).
Show Worked Solution

   \(\ce{HCl(aq) -> H+ (aq) + Cl– (aq)}\)

   \(\ce{C2H3COOH(aq) \rightleftharpoons H+ (aq) + C2H3COO– (aq)}\)

  • \(\ce{HCl}\) is a strong acid that fully dissociates in water, resulting in a high concentration of \(\ce{H+}\) ions and a low pH.
  • Acrylic acid, on the other hand, is a weak acid that only partially dissociates in water, resulting in a lower concentration of \(\ce{H+}\) ions and a higher pH.
  • When the concentration of \(\ce{HCl}\) decreases by a factor of 10, its pH increases by 1 due to the decrease in \(\ce{H+}\).
  • However, when the concentration of acrylic acid decreases by a factor of 10, its pH increases by less than 1.
  • This is due to the change in pH causing the equilibrium to shift right, producing more \(\ce{H+}\) ions in response to the dilution, resulting in a smaller change in the concentration of \(\ce{H+}\), and thus smaller change in pH.
  • At very dilute concentrations, the degree of dissociation of acrylic acid approaches 100% and its pH converges closely to that of \(\ce{HCl}\).

CHEMISTRY, M6 2022 HSC 25

The pH of two aqueous solutions was compared.
 

Explain why the `\text{HCN}(aq)` solution has a higher pH than the `\text{HCl}(aq)` solution. Include a relevant chemical equation for the `\text{HCN}(aq)` solution.   (3 marks)

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  • \( \ce{HCl} \) is a strong acid, ie it completely ionises in water to form \( \ce{H+} \) ions.
  • On the other hand, \( \ce{HCN} \) is a weak acid, ie it partially ionises in water to form \( \ce{H+} \) ions.
  •    \( \ce{HCl(aq) -> H+ (aq) + Cl- (aq)} \)
  •    \( \ce{HCN(aq) \rightleftharpoons H+ (aq) + CN– (aq)}\)
  • As \( \ce{[H+]} \) decreases, pH increases  (\( \ce{\text{pH} = – log [H+]} \))
  • Therefore, at the same 0.2M, the \( \ce{HCN} \) solution would have a lower \( \ce{[H+]} \) and thus would have a higher pH than \( \ce{HCl} \).
Show Worked Solution
  • \( \ce{HCl} \) is a strong acid, ie it completely ionises in water to form \( \ce{H+} \) ions.
  • On the other hand, \( \ce{HCN} \) is a weak acid, ie it partially ionises in water to form \( \ce{H+} \) ions.
  •    \( \ce{HCl(aq) -> H+ (aq) + Cl- (aq)} \)
  •    \( \ce{HCN(aq) \rightleftharpoons H+ (aq) + CN– (aq)}\)
  • As \( \ce{[H+]} \) decreases, pH increases  (\( \ce{\text{pH} = – log [H+]} \))
  • Therefore, at the same 0.2M, the \( \ce{HCN} \) solution would have a lower \( \ce{[H+]} \) and thus would have a higher pH than \( \ce{HCl} \).

Mean mark 57%.

CHEMISTRY, M6 2022 HSC 22

The following equation describes an equilibrium reaction.

\( \ce{HF(aq) + PO4^3-(aq) \rightleftharpoons HPO4^2-(aq) + F-(aq)} \)

Identify ONE base and its conjugate acid in the above equation.   (2 marks)
 

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Possible answers:

\begin{array} {ll}
\text{Base} & \text{Conjugate Acid} \\
\hline \ce{PO4^3-(aq)} & \ce{HPO4^2-(aq)} \\
\ce{F-(aq)} & \ce{HF(aq)} \\
  \end{array}

Show Worked Solution

Possible answers:

\begin{array} {ll}
\text{Base} & \text{Conjugate Acid} \\
\hline \ce{PO4^3-(aq)} & \ce{HPO4^2-(aq)} \\
\ce{F-(aq)} & \ce{HF(aq)} \\
  \end{array}