Reactivity 3.1.7 – Acid–Base Neutralization Reactions

Neutralization

• A neutralization reaction is a type of chemical reaction in which an acid reacts with a base to produce a salt and usually water.

• This process occurs through the combination of hydrogen ions (H⁺) from the acid and hydroxide ions (OH⁻) from the base, forming water:

\( \text{H}^+(aq) + \text{OH}^-(aq) \rightarrow \text{H}_2\text{O}(l) \)           

• In the Brønsted–Lowry acid–base model, the acid is defined as a proton donor and the base as a proton acceptor. Neutralization is thus a proton transfer reaction where the base receives a proton from the acid.
• For strong acids and strong bases: Both dissociate fully in solution, so the reaction proceeds completely and quickly, producing a neutral solution with pH = 7.
• For weak acids or weak bases: The reaction does not go to completion, and the pH of the final solution depends on the relative strengths of the acid and base involved.

• Salt formation: The salt produced in the reaction is composed of the cation from the base and the anion from the acid. For example, in the reaction between HCl and NaOH:

\( \text{HCl} + \text{NaOH} \rightarrow \text{NaCl} + \text{H}_2\text{O} \)

• Neutralization is usually exothermic, meaning it releases heat, which is why the reaction mixture often becomes warm.

• Not all neutralizations produce water: When weak bases like ammonia (NH₃) are used, the acid donates a proton to the base without forming water:

\( \text{NH}_3 + \text{HCl} \rightarrow \text{NH}_4^+ + \text{Cl}^- \)

Common Bases Involved in Neutralization:

• Metal hydroxides (e.g. NaOH, KOH)
• Metal oxides (e.g. CuO, MgO)
• Ammonia and amines (e.g. NH₃, CH₃NH₂)
• Soluble carbonates (e.g. Na₂CO₃, K₂CO₃)
• Hydrogencarbonates (e.g. NaHCO₃, KHCO₃)

Typical Ionic Equation (Strong Acid + Strong Base):

• \( \text{H}^+(aq) + \text{OH}^-(aq) \rightarrow \text{H}_2\text{O}(l) \)
• This is the net ionic equation for all strong acid–strong base reactions.

Energy Change in Neutralization:

• Standard enthalpy change of neutralization is typically around –57 kJ·mol⁻¹ for strong acid–strong base reactions.
• Values may vary for weak acids/bases due to partial ionization.

Key IB Concepts to Remember:

• Neutralization forms a salt and water only.
• The salt is formed from the cation of the base and the anion of the acid.
• pH at equivalence point:
  • Strong acid + strong base → pH = 7
  • Strong acid + weak base → pH < 7
  • Weak acid + strong base → pH > 7
• Do not confuse neutralization (a chemical process) with neutrality (pH = 7).

Reactions of Acids with Metal, Metal Oxides, Hydroxides, Carbonates, and Hydrogencarbonates

• Acids react with a range of bases to produce a salt and either water or carbon dioxide and water.
• This is an extension of neutralization reactions, where the base is not limited to hydroxide ions but may include metal oxides, carbonates, and hydrogen carbonates (bicarbonates).
• These are all proton transfer reactions that can be represented using molecular, ionic, and net ionic equations.

1. Acids Reacting with Metals

• When acids react with reactive metals, such as magnesium or zinc, a salt and hydrogen gas are produced.
• This is a redox reaction: the metal is oxidized (loses electrons), and H⁺ from the acid is reduced to H₂ gas.
• General word equation:

Metal + Acid → Salt + Hydrogen gas

• General ionic equation:

\( \text{M}(s) + 2\text{H}^+(aq) \rightarrow \text{M}^{2+}(aq) + \text{H}_2(g) \)

• Examples:
  • \( \text{Mg}(s) + 2\text{HCl}(aq) \rightarrow \text{MgCl}_2(aq) + \text{H}_2(g) \)

• • \( \text{Zn}(s) + \text{H}_2\text{SO}_4(aq) \rightarrow \text{ZnSO}_4(aq) + \text{H}_2(g) \)
• Only metals above hydrogen in the reactivity series react with acids to produce H₂ gas.

2. Reactions of Acids with Metal Oxides

• Metal oxides are generally basic in nature — many are insoluble bases.
• When a metal oxide reacts with an acid, a salt and water are formed. This is a classic neutralization reaction.
• General word equation:

Acid + Metal oxide → Salt + Water

• General ionic equation:

\( \text{2H}^+(aq) + \text{O}^{2-}(s) \rightarrow \text{H}_2\text{O}(l) \)

• The metal oxide does not dissolve entirely but still neutralizes the acid.
• Examples:
  • \( \text{H}_2\text{SO}_4(aq) + \text{CuO}(s) \rightarrow \text{CuSO}_4(aq) + \text{H}_2\text{O}(l) \)
  • \( 2\text{HCl}(aq) + \text{MgO}(s) \rightarrow \text{MgCl}_2(aq) + \text{H}_2\text{O}(l) \)
• This reaction is particularly useful in preparing salts in the lab (e.g. copper(II) sulfate).

3. Reactions of Acids with Metal Hydroxides

• Metal hydroxides are alkalis if soluble, and are classic bases.
• When acids react with metal hydroxides, a salt and water are produced — a classic neutralization reaction.
• General word equation:

Acid + Metal hydroxide → Salt + Water

• General ionic equation:

\( \text{H}^+(aq) + \text{OH}^-(aq) \rightarrow \text{H}_2\text{O}(l) \)

• This reaction is typically fast, especially with group 1 and group 2 hydroxides, which are highly soluble in water.
• Examples:
  • \( \text{HCl}(aq) + \text{NaOH}(aq) \rightarrow \text{NaCl}(aq) + \text{H}_2\text{O}(l) \)
  • \( 2\text{HNO}_3(aq) + \text{Ba(OH)}_2(aq) \rightarrow \text{Ba(NO}_3)_2(aq) + 2\text{H}_2\text{O}(l) \)
  • \( \text{H}_2\text{SO}_4(aq) + 2\text{KOH}(aq) \rightarrow \text{K}_2\text{SO}_4(aq) + 2\text{H}_2\text{O}(l) \)

• These reactions are often used in titrations to determine unknown concentrations of acids or bases.

4. Reactions of Acids with Metal Carbonates

• Metal carbonates react with acids to produce a salt, carbon dioxide gas, and water.
• This is an important reaction in identifying carbonates — the effervescence (bubbling) of CO₂ is a visual sign of reaction.
• General word equation:

Acid + Metal carbonate → Salt + Water + Carbon dioxide

• General ionic equation:

\( 2\text{H}^+(aq) + \text{CO}_3^{2-}(aq) \rightarrow \text{H}_2\text{O}(l) + \text{CO}_2(g) \)

• Always identify the parent acid and metal ion to determine the name of the salt formed.
• Examples:
  • \( \text{H}_2\text{SO}_4(aq) + \text{CaCO}_3(s) \rightarrow \text{CaSO}_4(aq) + \text{H}_2\text{O}(l) + \text{CO}_2(g) \)
  • \( 2\text{HCl}(aq) + \text{Na}_2\text{CO}_3(aq) \rightarrow 2\text{NaCl}(aq) + \text{H}_2\text{O}(l) + \text{CO}_2(g) \)
  • \( 2\text{CH}_3\text{COOH}(aq) + \text{MgCO}_3(s) \rightarrow (\text{CH}_3\text{COO})_2\text{Mg}(aq) + \text{H}_2\text{O}(l) + \text{CO}_2(g) \)

• This reaction helps in qualitative analysis — if a gas is produced that turns limewater milky, it confirms CO₂.

5. Reactions of Acids with Metal Hydrogencarbonates

• Metal hydrogencarbonates (also called metal bicarbonates) react with acids to produce a salt, water, and carbon dioxide gas.
• This is a classic <strongacid–base neutralization accompanied by decomposition of the hydrogencarbonate ion.
• General word equation:

Acid + Metal hydrogencarbonate → Salt + Water + Carbon dioxide

• General ionic equation:

\( \text{H}^+(aq) + \text{HCO}_3^-(aq) \rightarrow \text{H}_2\text{O}(l) + \text{CO}_2(g) \)

• The effervescence of CO₂ gas is a visual indicator of this reaction, just like with carbonates.

• Examples:
  • \( \text{HCl}(aq) + \text{NaHCO}_3(s) \rightarrow \text{NaCl}(aq) + \text{H}_2\text{O}(l) + \text{CO}_2(g) \)
  • \( 2\text{CH}_3\text{COOH}(aq) + \text{Ca(HCO}_3)_2(aq) \rightarrow (\text{CH}_3\text{COO})_2\text{Ca}(aq) + 2\text{H}_2\text{O}(l) + 2\text{CO}_2(g) \)
  • \( \text{H}_2\text{SO}_4(aq) + 2\text{KHCO}_3(aq) \rightarrow \text{K}_2\text{SO}_4(aq) + 2\text{H}_2\text{O}(l) + 2\text{CO}_2(g) \)
• These reactions are often used in medicine (antacids) and baking (baking soda + acid).

General Reaction Types

Reactions with Ammonia and Amines

• Ammonia and amines are weak bases that react with acids to form ammonium salts.
• No water is formed because there are no hydroxide ions present to combine with hydrogen ions.

Important Concepts

• All these reactions are acid–base reactions under the Brønsted–Lowry theory (proton transfer).
• Metal oxides and hydroxides act as bases by accepting H⁺ to form water.
• Carbonates and hydrogencarbonates release CO₂ upon reacting with acids.
• When formulating equations, always ensure charge balance and state symbols are included.