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CHEMISTRY, M6 2025 HSC 11 MC

The structures of two substances, \(\text{X}\) and \(\text{Y}\), are shown.
 

Which row of the table correctly classifies these substances as a Brønsted-Lowry acid or a Brønsted-Lowry base?
 

\begin{align*}
\begin{array}{l}
\rule{0pt}{2.5ex} \ & \\
 \ \rule[-1ex]{0pt}{0pt}& \\
\rule{0pt}{2.5ex}\textbf{A.}\rule[-1ex]{0pt}{0pt}\\
\rule{0pt}{2.5ex}\textbf{B.}\rule[-1ex]{0pt}{0pt}\\
\rule{0pt}{2.5ex}\textbf{C.}\rule[-1ex]{0pt}{0pt}\\
\rule{0pt}{2.5ex}\textbf{D.}\rule[-1ex]{0pt}{0pt}\\
\end{array}
\begin{array}{|c|c|}
\hline
\rule{0pt}{2.5ex}\textit{Brønsted-Lowry} & \textit{Brønsted-Lowry} \\
\textit{acid}\rule[-1ex]{0pt}{0pt}& \textit{base} \\
\hline
\rule{0pt}{2.5ex}\text{-}\rule[-1ex]{0pt}{0pt}&\text{X and Y}\\
\hline
\rule{0pt}{2.5ex}\text{X and Y}\rule[-1ex]{0pt}{0pt}& \text{-}\\
\hline
\rule{0pt}{2.5ex}\text{Y}\rule[-1ex]{0pt}{0pt}& \text{X} \\
\hline
\rule{0pt}{2.5ex}\text{X}\rule[-1ex]{0pt}{0pt}& \text{Y} \\
\hline
\end{array}
\end{align*}

Show Answers Only

\(A\)

Show Worked Solution
  • A Brønsted-Lowry acid is a proton donor and a Brønsted-Lowry base is a proton acceptor.
  • \(X\) is propanoate (the conjugate base of propanoic acid) and is therefore a proton accepter making it a Brønsted-Lowry base.
  • \(Y\) is ethanamine and is considered a weak base where the \(\ce{NH2}\) group can accept a proton to become \(\ce{NH3+}\).

\(\Rightarrow A\)

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)

    •  \( \ce{HCl(aq)} \)
    • \( \ce{NH4Cl(aq)} \)
Show Answers Only
  • 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)

Show Answers Only

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)

Show Answers Only
  • 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)

Show Answers Only

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

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

  • \(\ce{HCl}\) 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)

Show Answers Only
  • \( \ce{HCl} \) 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}