1.2 Types Of Chemical Reactions

Chemical reactions involve the breaking of old chemical bonds within reactant molecules and the formation of new chemical bonds to create product molecules. Atoms are simply rearranged, not created or destroyed.

1.2.1 Combination Reaction

In a combination reaction, two or more reactants (elements or compounds) combine to form a single product.

Example: When calcium oxide (quicklime, $\text{CaO}$) reacts with water ($\text{H}_2\text{O}$), it forms calcium hydroxide (slaked lime, $\text{Ca(OH)}_2$). This reaction is vigorous and releases a large amount of heat.

$\text{CaO(s)} + \text{H}_2\text{O(l)} \longrightarrow \text{Ca(OH)}_2\text{(aq)} + \text{Heat}$

Here, two substances ($\text{CaO}$ and $\text{H}_2\text{O}$) combine to form one substance ($\text{Ca(OH)}_2$), fitting the definition of a combination reaction.

This reaction is also an example of an exothermic reaction because heat is released during the process, causing the reaction mixture to become warm.

Other Examples of Combination Reactions:

• Burning of carbon (coal): $\text{C(s)} + \text{O}_2\text{(g)} \longrightarrow \text{CO}_2\text{(g)}$
• Formation of water: $2\text{H}_2\text{(g)} + \text{O}_2\text{(g)} \longrightarrow 2\text{H}_2\text{O(l)}$

Exothermic Reactions: Reactions that release energy (usually as heat). Examples include:

• Burning of natural gas (methane): $\text{CH}_4\text{(g)} + 2\text{O}_2\text{(g)} \longrightarrow \text{CO}_2\text{(g)} + 2\text{H}_2\text{O(g)} + \text{Heat}$
• Respiration: The process in living cells where glucose combines with oxygen to produce carbon dioxide, water, and energy. $\text{C}_6\text{H}_{12}\text{O}_6\text{(aq)} + 6\text{O}_2\text{(aq)} \longrightarrow 6\text{CO}_2\text{(aq)} + 6\text{H}_2\text{O(l)} + \text{energy}$
• Decomposition of vegetable matter into compost.

Calcium hydroxide solution (slaked lime) is used for whitewashing walls. It slowly reacts with carbon dioxide in the air to form a thin, shiny layer of calcium carbonate ($\text{CaCO}_3$) after two to three days. $\text{Ca(OH)}_2\text{(aq)} + \text{CO}_2\text{(g)} \longrightarrow \text{CaCO}_3\text{(s)} + \text{H}_2\text{O(l)}$

1.2.2 Decomposition Reaction

A decomposition reaction is the opposite of a combination reaction. In a decomposition reaction, a single reactant breaks down to give two or more simpler products.

Decomposition reactions typically require energy input to break the bonds in the reactant. This energy can be supplied in the form of heat, light, or electricity.

• Thermal Decomposition: Decomposition caused by heating.

  Example 1: Heating ferrous sulphate crystals ($\text{FeSO}_4 \cdot 7\text{H}_2\text{O}$, green colour). When heated, they first lose water and the colour changes. Then, they decompose into ferric oxide ($\text{Fe}_2\text{O}_3$, brown/black solid), sulphur dioxide ($\text{SO}_2$, gas with characteristic burning sulphur smell), and sulphur trioxide ($\text{SO}_3$, gas).

  $2\text{FeSO}_4\text{(s)} \xrightarrow{\text{Heat}} \text{Fe}_2\text{O}_3\text{(s)} + \text{SO}_2\text{(g)} + \text{SO}_3\text{(g)}$

  Example 2: Heating calcium carbonate ($\text{CaCO}_3$, limestone). It decomposes into calcium oxide (quicklime, $\text{CaO}$) and carbon dioxide ($\text{CO}_2$). This is an important industrial process for manufacturing cement.

  $\text{CaCO}_3\text{(s)} \xrightarrow{\text{Heat}} \text{CaO(s)} + \text{CO}_2\text{(g)}$

  Example 3: Heating lead nitrate ($\text{Pb(NO}_3)_2$, white powder). It decomposes to lead oxide ($\text{PbO}$, yellow solid), nitrogen dioxide ($\text{NO}_2$, brown fumes), and oxygen gas ($\text{O}_2$).

  $2\text{Pb(NO}_3)_2\text{(s)} \xrightarrow{\text{Heat}} 2\text{PbO(s)} + 4\text{NO}_2\text{(g)} + \text{O}_2\text{(g)}$

• Electrolytic Decomposition (Electrolysis): Decomposition caused by passing electric current.

  Example: Electrolysis of water ($\text{H}_2\text{O}$). Water decomposes into hydrogen gas ($\text{H}_2$) and oxygen gas ($\text{O}_2$) when electric current is passed through it (often with a little acid added to make it conduct electricity). The volume of hydrogen gas collected at one electrode is twice the volume of oxygen gas collected at the other, consistent with the $2:1$ ratio of hydrogen to oxygen atoms in water.

  $2\text{H}_2\text{O(l)} \xrightarrow{\text{Electricity}} 2\text{H}_2\text{(g)} + \text{O}_2\text{(g)}$

• Photochemical Decomposition (Photolysis): Decomposition caused by light energy.

  Example 1: Exposing silver chloride ($\text{AgCl}$, white solid) to sunlight. It decomposes into silver metal ($\text{Ag}$, grey solid) and chlorine gas ($\text{Cl}_2$).

  $2\text{AgCl(s)} \xrightarrow{\text{Sunlight}} 2\text{Ag(s)} + \text{Cl}_2\text{(g)}$

  Example 2: Silver bromide ($\text{AgBr}$) also undergoes photochemical decomposition in sunlight, forming silver and bromine gas. These reactions are used in traditional black and white photography.

  $2\text{AgBr(s)} \xrightarrow{\text{Sunlight}} 2\text{Ag(s)} + \text{Br}_2\text{(g)}$

Reactions that absorb energy are called endothermic reactions. Decomposition reactions are typically endothermic because they require energy input to break chemical bonds.

Example of an Endothermic Reaction (non-decomposition): Mixing barium hydroxide and ammonium chloride feels cold because the reaction absorbs heat from the surroundings.

1.2.3 Displacement Reaction

A displacement reaction is a type of chemical reaction where a more reactive element displaces (removes) a less reactive element from its compound.

Example: When an iron nail ($\text{Fe}$) is dipped into a copper sulphate solution ($\text{CuSO}_4$, blue colour), iron displaces copper from the copper sulphate solution. The iron nail becomes coated with brownish copper metal, and the blue colour of the copper sulphate solution fades due to the formation of iron sulphate ($\text{FeSO}_4$, which is typically green or colourless in dilute solution).

$\text{Fe(s)} + \text{CuSO}_4\text{(aq)} \longrightarrow \text{FeSO}_4\text{(aq)} + \text{Cu(s)}$

Here, iron is more reactive than copper, so it displaces copper from the copper sulphate solution.

Other Examples of Displacement Reactions:

• Zinc ($\text{Zn}$) reacting with copper sulphate ($\text{CuSO}_4$): Zinc is more reactive than copper.

  $\text{Zn(s)} + \text{CuSO}_4\text{(aq)} \longrightarrow \text{ZnSO}_4\text{(aq)} + \text{Cu(s)}$

• Lead ($\text{Pb}$) reacting with copper chloride ($\text{CuCl}_2$): Lead is more reactive than copper.

  $\text{Pb(s)} + \text{CuCl}_2\text{(aq)} \longrightarrow \text{PbCl}_2\text{(aq)} + \text{Cu(s)}$

The relative reactivity of elements determines which element can displace another. More reactive metals are generally higher in the reactivity series.

1.2.4 Double Displacement Reaction

In a double displacement reaction, there is an exchange of ions between two reactant compounds, leading to the formation of two new compounds.

These reactions often occur in aqueous solutions and can result in the formation of a precipitate, a gas, or water.

Example: Mixing a solution of sodium sulphate ($\text{Na}_2\text{SO}_4$) with a solution of barium chloride ($\text{BaCl}_2$). A white insoluble substance, barium sulphate ($\text{BaSO}_4$), is formed as a precipitate. Sodium chloride ($\text{NaCl}$) is also formed and remains dissolved in the solution.

$\text{Na}_2\text{SO}_4\text{(aq)} + \text{BaCl}_2\text{(aq)} \longrightarrow \text{BaSO}_4\text{(s)} \downarrow + 2\text{NaCl(aq)}$

Here, the sodium ions ($\text{Na}^+$) from $\text{Na}_2\text{SO}_4$ exchange partners with the barium ions ($\text{Ba}^{2+}$) from $\text{BaCl}_2$. Sodium combines with chloride ($\text{Cl}^-$) to form $\text{NaCl}$, and barium combines with sulphate ($\text{SO}_4^{2-}$) to form $\text{BaSO}_4$.

Reactions that produce a precipitate are also called precipitation reactions.

Example: Mixing lead(II) nitrate solution ($\text{Pb(NO}_3)_2$) and potassium iodide solution ($\text{KI}$) forms a yellow precipitate of lead(II) iodide ($\text{PbI}_2$) and potassium nitrate ($\text{KNO}_3$). This is a double displacement reaction.

$\text{Pb(NO}_3)_2\text{(aq)} + 2\text{KI(aq)} \longrightarrow \text{PbI}_2\text{(s)} \downarrow + 2\text{KNO}_3\text{(aq)}$

1.2.5 Oxidation And Reduction

Oxidation and reduction are fundamental concepts in chemistry that describe the transfer or sharing of electrons or the gain/loss of oxygen and hydrogen during chemical reactions.

Originally, these terms were defined based on the gain or loss of oxygen or hydrogen:

• Oxidation: The process involving the gain of oxygen or the loss of hydrogen by a substance during a reaction.
• Reduction: The process involving the loss of oxygen or the gain of hydrogen by a substance during a reaction.

Example: Heating copper powder ($\text{Cu}$, shiny brown) in air (oxygen, $\text{O}_2$). The copper reacts with oxygen to form black copper(II) oxide ($\text{CuO}$).

$2\text{Cu(s)} + \text{O}_2\text{(g)} \xrightarrow{\text{Heat}} 2\text{CuO(s)}$

Here, copper ($\text{Cu}$) gains oxygen, so it is oxidised to copper(II) oxide ($\text{CuO}$).

If hydrogen gas ($\text{H}_2$) is passed over the heated copper(II) oxide ($\text{CuO}$), the black coating turns brown as copper is recovered, and water ($\text{H}_2\text{O}$) is formed.

$\text{CuO(s)} + \text{H}_2\text{(g)} \xrightarrow{\text{Heat}} \text{Cu(s)} + \text{H}_2\text{O(l)}$

In this reaction:

• Copper(II) oxide ($\text{CuO}$) loses oxygen, so it is reduced to copper ($\text{Cu}$).
• Hydrogen ($\text{H}_2$) gains oxygen, so it is oxidised to water ($\text{H}_2\text{O}$).

Reactions where one substance is oxidised and another substance is reduced simultaneously are called oxidation-reduction reactions or redox reactions. Oxidation and reduction always occur together; one cannot happen without the other.

Other Examples of Redox Reactions:

• $\text{ZnO + C} \longrightarrow \text{Zn + CO}$ (ZnO is reduced, C is oxidised)
• $\text{MnO}_2 + 4\text{HCl} \longrightarrow \text{MnCl}_2 + 2\text{H}_2\text{O} + \text{Cl}_2$ (MnO₂ is reduced, HCl is oxidised)

A substance that causes oxidation (by providing oxygen or removing hydrogen) is called an oxidising agent. A substance that causes reduction (by providing hydrogen or removing oxygen) is called a reducing agent. In the example $\text{CuO} + \text{H}_2 \longrightarrow \text{Cu} + \text{H}_2\text{O}$, $\text{CuO}$ is the oxidising agent, and $\text{H}_2$ is the reducing agent.

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