Reactions of Acids Revision Notes for HSC SSCE Chemistry

Reactions of Acids

Introduction

Acids can be identified not only by their properties but also by the characteristic types of reactions they undergo. Because acids share similar properties, they follow predictable patterns in their chemical reactions. Understanding these patterns allows us to predict the products formed when acids react with different substances.

infoNote

This note covers three main types of acid reactions:

Reactions with bases (neutralisation)
Reactions with carbonates
Reactions with metals

By mastering these three reaction patterns, you'll be able to predict the products of most acid reactions you encounter in chemistry.

Acid-base reactions

What is neutralisation?

Neutralisation is the chemical reaction between an acid and a base that produces a salt and water. When the correct amounts of acid and base react together, the resulting solution is neutral—meaning it is neither acidic nor basic.

The process of neutralisation follows a general pattern that can be expressed as:

$\text{Acid} + \text{base} \rightarrow \text{salt} + \text{water}$

This general equation is incredibly useful because it allows us to predict the products of most acid-base reactions.

Types of acid-base reactions

Reaction with hydroxides

When an acid reacts with a hydroxide compound, neutralisation occurs to produce a salt and water:

$\text{Acid} + \text{hydroxide} \rightarrow \text{salt} + \text{water}$

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Worked Example: Hydrochloric acid reacting with sodium hydroxide

The complete chemical equation:

$\text{HCl}_{(aq)} + \text{NaOH}_{(aq)} \rightarrow \text{NaCl}_{(aq)} + \text{H}_2\text{O}_{(l)}$

In this reaction, sodium chloride (common salt) and water are produced. The sodium ions ( $\text{Na}^+$ ) and chloride ions ( $\text{Cl}^-$ ) are spectator ions—they don't actually participate in the reaction.

The actual neutralisation involves hydrogen ions from the acid combining with hydroxide ions from the base:

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

chatImportant

This is called the net ionic equation and shows the essence of all acid-hydroxide neutralisation reactions. This equation is fundamental to understanding neutralisation—it's the same for every acid-hydroxide reaction, regardless of which specific acid or base you use.

Reaction with basic or amphoteric oxides

Acids also undergo neutralisation reactions with basic or amphoteric oxides:

$\text{Acid} + \text{oxide} \rightarrow \text{salt} + \text{water}$

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Worked Example: Sulfuric acid reacting with magnesium oxide

$\text{H}_2\text{SO}_{4(aq)} + \text{MgO}_{(s)} \rightarrow \text{MgSO}_{4(aq)} + \text{H}_2\text{O}_{(l)}$

This reaction produces magnesium sulfate (a salt) and water. Basic oxides like magnesium oxide act as bases when they react with acids.

Special case: reactions with ammonia

The reaction between acids and ammonia ( $\text{NH}_3$ ) is an important exception to the general neutralisation pattern. Unlike typical acid-base reactions, this one does not produce water in the same way.

When hydrochloric acid reacts with ammonia:

$\text{HCl}_{(aq)} + \text{NH}_{3(aq)} \rightarrow \text{NH}_4\text{Cl}_{(aq)}$

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In this reaction, the hydrogen ion ( $\text{H}^+$ ) from the acid attaches directly to the ammonia molecule to form the ammonium ion ( $\text{NH}_4^+$ ), rather than combining with hydroxide to form water.

Ammonia is considered a weak base because when it dissolves in water, it produces hydroxide ions:

\text{NH}_3(g) + \text{H}_2\text{O}(l) \rightarrow \text{NH}_4^+(aq) + \text{OH}^-(aq)

The ammonium ion ( $\text{NH}_4^+$ ) formed is itself a weak acid. This means when ammonium salts react with bases, they release ammonia gas:

$\text{NaOH}_{(aq)} + \text{NH}_4\text{Cl}_{(aq)} \rightarrow \text{NaCl}_{(aq)} + \text{NH}_{3(g)} + \text{H}_2\text{O}_{(l)}$

Real-world application: fertilisers

Ammonium salts like ammonium sulfate and ammonium nitrate are commonly used as fertilisers.

chatImportant

Farmers must be careful not to apply ammonium compounds at the same time as alkaline substances like lime. If they do, the ammonium salt reacts with the alkali to release ammonia gas into the atmosphere, meaning the nitrogen is lost and unavailable to plants.

Acid-carbonate reactions

Formation of carbon dioxide

When acids react with carbonates, they produce three products: a salt, carbon dioxide gas, and water. This reaction is responsible for the characteristic fizzing and bubbling observed when acids meet carbonates.

The general equation for this reaction is:

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

Reactions with carbonates

A carbonate ion has the formula $\text{CO}_3^{2-}$ . Common carbonates include calcium carbonate (found in limestone and marble) and sodium carbonate.

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Worked Example 1: Hydrochloric acid reacting with calcium carbonate

$2\text{HCl}_{(aq)} + \text{CaCO}_{3(s)} \rightarrow \text{CaCl}_{2(aq)} + \text{H}_2\text{O}_{(l)} + \text{CO}_{2(g)}$

Worked Example 2: Sulfuric acid reacting with calcium carbonate

$\text{H}_2\text{SO}_{4(aq)} + \text{CaCO}_{3(s)} \rightarrow \text{CaSO}_{4(aq)} + \text{H}_2\text{O}_{(l)} + \text{CO}_{2(g)}$

These reactions are vigorous, producing lots of fizzing as carbon dioxide gas escapes from the solution.

Reactions with hydrogen carbonates

A hydrogen carbonate ion has the formula $\text{HCO}_3^-$ . These compounds undergo very similar reactions to carbonates, also producing a salt, water, and carbon dioxide:

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

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Worked Example 1: Hydrochloric acid reacting with sodium hydrogen carbonate

$\text{HCl}_{(aq)} + \text{NaHCO}_{3(s)} \rightarrow \text{NaCl}_{(aq)} + \text{H}_2\text{O}_{(l)} + \text{CO}_{2(g)}$

Worked Example 2: Sulfuric acid reacting with sodium hydrogen carbonate

$\text{H}_2\text{SO}_{4(aq)} + 2\text{NaHCO}_{3(s)} \rightarrow \text{Na}_2\text{SO}_{4(aq)} + 2\text{H}_2\text{O}_{(l)} + 2\text{CO}_{2(g)}$

Real-world application: baking

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Practical Application in the Kitchen

Acid-carbonate reactions are essential in cooking. When baking bread or sponge cakes, the carbon dioxide gas produced by these reactions creates bubbles in the mixture, causing it to rise and giving baked goods their characteristic honeycomb structure.

Baking powder typically contains a carbonate or hydrogen carbonate along with an acid, which react together when water is added and the mixture is heated.

Acid-metal reactions

Formation of hydrogen gas

Most metals react with dilute acids to produce a salt and hydrogen gas. This is the same type of reaction that occurs when acid rain corrodes metal structures and vehicles.

The general equation for acid-metal reactions is:

$\text{Dilute acid} + \text{metal} \rightarrow \text{salt} + \text{hydrogen gas}$

Which metals react?

Most metals react with dilute hydrochloric acid and sulfuric acid, including:

Magnesium
Zinc
Iron
Aluminium
Calcium

chatImportant

Some metals do not react with dilute acids:

Copper
Silver
Gold
Platinum

Tin and lead do react, but very slowly. This is why copper is often used for pipes and containers that might come into contact with acidic substances.

Examples of acid-metal reactions

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Worked Example 1: Hydrochloric acid reacting with magnesium metal

$2\text{HCl}_{(aq)} + \text{Mg}_{(s)} \rightarrow \text{MgCl}_{2(aq)} + \text{H}_{2(g)}$

Worked Example 2: Sulfuric acid reacting with magnesium metal

$\text{H}_2\text{SO}_{4(aq)} + \text{Mg}_{(s)} \rightarrow \text{MgSO}_{4(aq)} + \text{H}_{2(g)}$

These reactions produce visible hydrogen gas bubbles and often generate heat.

Net ionic equation

It's important to understand that the hydrogen ions ( $\text{H}^+$ ) from the acid are what actually react with the metal. This can be shown in a net ionic equation:

$2\text{H}^+_{(aq)} + \text{Mg}_{(s)} \rightarrow \text{Mg}^{2+}_{(aq)} + \text{H}_{2(g)}$

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This equation shows that hydrogen ions are reduced to hydrogen gas, while the metal is oxidised to form metal ions in solution. This is a redox reaction—reduction and oxidation occurring simultaneously.

Real-world application: acid rain and corrosion

Acid rain forms when pollutants in the atmosphere dissolve in rainwater, making it acidic. This acidic rain reacts with metal structures, causing them to corrode much faster than they would in normal rain. Understanding acid-metal reactions helps us protect metal structures through appropriate coatings and materials selection.

Exam tips

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Key Strategies for Success:

When answering questions about acid reactions:

Identify the reactant type: Is it a base, carbonate, or metal?
Apply the correct general equation: Each reaction type has its own pattern
Balance your equations: Make sure atoms are conserved on both sides
Include state symbols: (s) for solid, (l) for liquid, (aq) for aqueous, (g) for gas
Look for keywords: "Fizzing" or "bubbles" suggests carbon dioxide or hydrogen gas production
Check for special cases: Remember ammonia reactions don't follow the typical pattern

Remember!

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Key Points to Remember:

Neutralisation is when acids and bases react to form a salt and water:

$\text{Acid} + \text{base} \rightarrow \text{salt} + \text{water}$

Acid-carbonate reactions produce three products: a salt, carbon dioxide gas, and water. This reaction causes fizzing and is used in baking.
Acid-metal reactions produce a salt and hydrogen gas. Most metals react with dilute acids, but copper, silver, gold, and platinum do not.
Pattern recognition is key: By learning these three general equations, you can predict the products of most acid reactions you'll encounter.
The hydrogen ion ( $\text{H}^+$ ) is the reactive species in all acid reactions—it's what makes acids acidic and determines how they react with other substances.

Reactions of Acids (HSC SSCE Chemistry): Revision Notes