Reactions of Acids and Bases (VCE SSCE Chemistry): Revision Notes

Reactions of Acids and Bases

Acids and bases were originally grouped together because of their similar chemical behaviour. They react readily with many different chemicals, and understanding these reaction patterns helps us predict what products will form.

In this topic, you will learn about three main types of acid-base reactions and how to represent them using ionic equations.

General reactions of acids and bases

While acids and bases can react in many ways, we can categorise their reactions based on what they react with and what products form. The three main reaction types you need to know are reactions of acids with:

• metal hydroxides
• metal carbonates and hydrogen carbonates
• reactive metals

Acids and metal hydroxides

When acids react with metal hydroxides, a neutralisation reaction occurs. These reactions produce two products: a salt and water.

Understanding neutralisation

Soluble metal hydroxides (such as $\text{NaOH}$, $\text{Ca(OH)}_2$, and $\text{Mg(OH)}_2$) dissociate in water to form metal cations and hydroxide ions ($\text{OH}^-$(aq)). The hydroxide ions from the base react with the hydrogen ions ($\text{H}^+$(aq)) from the acid to form water molecules.

The general equation for this reaction is:

$\text{acid} + \text{metal hydroxide} \rightarrow \text{salt} + \text{water}$

What is a salt?

A salt is an ionic compound that consists of:

• The positive ion (cation) from the base
• The negative ion (anion) from the acid

Common salts from neutralisation reactions

The table below shows some common neutralisation reactions and the salts they produce:

Reactants (acid + metal hydroxide)	Name of salt formed	Formulas of ions present in the salt solution
hydrochloric acid + potassium hydroxide	potassium chloride	$\text{K}^+\text{(aq)} + \text{Cl}^-\text{(aq)}$
hydrochloric acid + magnesium hydroxide	magnesium chloride	$\text{Mg}^{2+}\text{(aq)} + \text{Cl}^-\text{(aq)}$
nitric acid + sodium hydroxide	sodium nitrate	$\text{Na}^+\text{(aq)} + \text{NO}_3^-\text{(aq)}$
sulfuric acid + zinc hydroxide	zinc sulfate	$\text{Zn}^{2+}\text{(aq)} + \text{SO}_4^{2-}\text{(aq)}$
phosphoric acid + potassium hydroxide	potassium phosphate	$\text{K}^+\text{(aq)} + \text{PO}_4^{3-}\text{(aq)}$
ethanoic acid + calcium hydroxide	calcium ethanoate	$\text{Ca}^{2+}\text{(aq)} + \text{CH}_3\text{COO}^-\text{(aq)}$

Representing neutralisation with ionic equations

When we write ionic equations, we show only the ions that actually participate in the reaction. Ions that remain unchanged are called spectator ions and are omitted from the ionic equation.

The diagram below shows what happens at the ionic level when hydrochloric acid and sodium hydroxide solutions are mixed:

Writing ionic equations step-by-step

Here's how to write an ionic equation for the reaction between hydrochloric acid and sodium hydroxide:

1. General reaction

Thinking: What is the general reaction? Identify the products formed.

Working: acid + metal hydroxide → salt + water

A solution of sodium chloride and water is formed.

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2. Identify reactants and products (with states)

Thinking: Identify the reactants and products. Indicate the state of each, i.e. (aq), (s), (l) or (g).

Working:

Reactants:

• HCl(aq) is ionised in solution, forming H⁺(aq) and Cl⁻(aq)
• NaOH(aq) is dissociated in solution, forming Na⁺(aq) and OH⁻(aq)

Products:

• Sodium chloride is dissociated and exists as Na⁺(aq) and Cl⁻(aq)
• Water is a molecular compound and its formula is H₂O(l)

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3. Write full ionic equation

Thinking: Write the equation showing all reactants and products, in ionised form where possible. (There is no need to balance the equation yet.)

Working:

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

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4. Identify spectator ions

Thinking: Identify the spectator ions: the ions that have an (aq) state, both as a reactant and as a product.

Working: Na⁺(aq) and Cl⁻(aq)

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5. Write net ionic equation

Thinking: Rewrite the equation without the spectator ions. Balance the equation with respect to atoms and charge.

Working:

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

Note: If hydronium ions are used instead of H⁺:

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

The final ionic equation is:

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

Alternatively, if we represent hydrogen ions as hydronium ions:

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

Practical applications of neutralisation

Neutralisation reactions are used in everyday situations to reduce the effects of acids or bases.

Treating insect bites and stings:

Methanoic acid ($\text{HCOOH}$) is released from the stings of ants, bees and nettles. If affected skin is rinsed with a dilute solution of ammonia (a weak base), the acid becomes neutralised and is no longer painful. In contrast, wasp venom is alkaline, so it can be neutralised by rinsing with vinegar (ethanoic acid).

Toothpaste:

Bacteria on tooth enamel feed on sugars in food and produce acids as waste products. These acids attack the enamel, leading to tooth decay. Toothpastes are formulated as weak bases to neutralise these food acids and protect teeth.

Acids and metal carbonates

When acids react with metal carbonates or metal hydrogen carbonates, they produce three products: a salt, water, and carbon dioxide gas. This type of reaction is responsible for the weathering of limestone buildings and statues.

Reactions with metal carbonates

Metal carbonates include compounds like sodium carbonate ($\text{Na}_2\text{CO}_3$), magnesium carbonate ($\text{MgCO}_3$), and calcium carbonate ($\text{CaCO}_3$).

The general equation for these reactions is:

$\text{acid} + \text{metal carbonate} \rightarrow \text{salt} + \text{water} + \text{carbon dioxide}$

The diagram below shows this reaction at the ionic level:

Reactions with metal hydrogen carbonates

Metal hydrogen carbonates (also called bicarbonates) include sodium hydrogen carbonate ($\text{NaHCO}_3$), potassium hydrogen carbonate ($\text{KHCO}_3$), and calcium hydrogen carbonate ($\text{Ca(HCO}_3)_2$). These also react with acids to produce carbon dioxide, salt, and water.

The general equation is:

$\text{acid} + \text{metal hydrogen carbonate} \rightarrow \text{salt} + \text{water} + \text{carbon dioxide}$

For example:

$\text{HCl(aq)} + \text{NaHCO}_3\text{(s)} \rightarrow \text{NaCl(aq)} + \text{H}_2\text{O(l)} + \text{CO}_2\text{(g)}$

Writing ionic equations for carbonate reactions

The process for writing ionic equations is the same as for neutralisation reactions. Here's an example for the reaction between nitric acid and magnesium carbonate:

1. General reaction

Thinking: What is the general reaction? Identify the products.

Working: acid + metal carbonate → salt + water + carbon dioxide

Products of this reaction are magnesium nitrate in solution, water and carbon dioxide gas.

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2. Identify reactants and products (with states)

Thinking: Identify the reactants and products. Indicate the state of each, i.e. (aq), (s), (l) or (g).

Working:

Reactants:

• Nitric acid is ionised in solution, forming H⁺(aq) and NO₃⁻(aq) ions
• Magnesium carbonate is an ionic solid, MgCO₃(s)

Products:

• Magnesium nitrate is dissociated into Mg²⁺(aq) and NO₃⁻(aq) ions
• Water has the formula H₂O(l)
• Carbon dioxide has the formula CO₂(g)

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3. Write full ionic equation

Thinking: Write the equation showing all reactants and products. (There is no need to balance the equation yet.)

Working:

$$\text{H}^+(aq) + \text{NO}_3^-(aq) + \text{MgCO}_3(s) \rightarrow \text{Mg}^{2+}(aq) + \text{NO}_3^-(aq) + \text{H}_2\text{O}(l) + \text{CO}_2(g)$$

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4. Identify spectator ions

Thinking: Identify the spectator ions.

Working: NO₃⁻(aq)

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5. Write net ionic equation and balance

Thinking: Rewrite the equation without the spectator ions. Balance the equation with respect to atoms and charge.

Working:

Unbalanced:

$$\text{H}^+(aq) + \text{MgCO}_3(s) \rightarrow \text{Mg}^{2+}(aq) + \text{H}_2\text{O}(l) + \text{CO}_2(g)$$

Balanced:

$$2\text{H}^+(aq) + \text{MgCO}_3(s) \rightarrow \text{Mg}^{2+}(aq) + \text{H}_2\text{O}(l) + \text{CO}_2(g)$$

Note: If hydronium ions are used instead of H⁺:

$$2\text{H}_3\text{O}^+(aq) + \text{MgCO}_3(s) \rightarrow \text{Mg}^{2+}(aq) + 3\text{H}_2\text{O}(l) + \text{CO}_2(g)$$

The final balanced ionic equation is:

$2\text{H}^+\text{(aq)} + \text{MgCO}_3\text{(s)} \rightarrow \text{Mg}^{2+}\text{(aq)} + \text{H}_2\text{O(l)} + \text{CO}_2\text{(g)}$

Or using hydronium ions:

$2\text{H}_3\text{O}^+\text{(aq)} + \text{MgCO}_3\text{(s)} \rightarrow \text{Mg}^{2+}\text{(aq)} + 3\text{H}_2\text{O(l)} + \text{CO}_2\text{(g)}$

Practical applications

Bicarbonate of soda in baking:

Self-raising flour contains tartaric acid and sodium hydrogen carbonate (bicarbonate of soda). When heated in the oven, these react to produce carbon dioxide gas, which causes cakes and muffins to rise.

Antacids:

Excess hydrochloric acid in the stomach can cause indigestion, heartburn, and upset stomach. Antacids contain bases that neutralise this excess acid. Common antacids include:

• Calcium carbonate and sodium hydrogen carbonate (quick-acting)
• Aluminium hydroxide and magnesium hydroxide (slower acting)

Testing for carbonate salts:

You can detect carbonate salts by adding an acid to an unknown sample. If carbon dioxide is produced, the sample contains carbonate.

The limewater test confirms the presence of carbon dioxide. Limewater is a saturated solution of calcium hydroxide ($\text{Ca(OH)}_2$). When carbon dioxide is bubbled through limewater, it turns milky or cloudy due to the precipitation of calcium carbonate:

$\text{Ca(OH)}_2\text{(aq)} + \text{CO}_2\text{(g)} \rightarrow \text{CaCO}_3\text{(s)} + \text{H}_2\text{O(l)}$

Acids and reactive metals

When dilute acids are added to reactive metals, bubbles of hydrogen gas are released and a salt is formed.

Understanding the reaction

The general equation for this reaction is:

$\text{acid} + \text{reactive metal} \rightarrow \text{salt} + \text{hydrogen}$

Reactive metals include calcium, magnesium, iron, and zinc. Copper, silver, and gold do not react with dilute acids.

Writing ionic equations for metal reactions

Here's the step-by-step process for writing the ionic equation for hydrochloric acid reacting with zinc:

1. General reaction

Thinking:
What is the general reaction? Identify the products formed.

Working:
acid + reactive metal → salt + hydrogen

Hydrogen gas and zinc chloride solution are produced.

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2. Identify reactants and products (with states)

Thinking:
Identify the reactants and products. Indicate the state of each, i.e. (aq), (s), (l) or (g).

Working:

Reactants:

• Zinc is a solid, Zn(s)
• Hydrochloric acid is ionised, forming H⁺(aq) and Cl⁻(aq) ions

Products:

• Hydrogen gas, H₂(g)
• Zinc chloride is dissociated into Zn²⁺(aq) and Cl⁻(aq) ions

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3. Write full ionic equation

Thinking:
Write the equation showing all reactants and products.
(There is no need to balance the equation yet.)

Working:

$$\text{H}^+(aq) + \text{Cl}^-(aq) + \text{Zn}(s) \rightarrow \text{Zn}^{2+}(aq) + \text{Cl}^-(aq) + \text{H}_2(g)$$

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4. Identify spectator ions

Thinking:
Identify the spectator ions.

Working:
Cl⁻(aq)

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5. Write net ionic equation and balance

Thinking:
Rewrite the equation without the spectator ions.

Working:

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

The final balanced ionic equation is:

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

Remember!