Science - Combined (0653) | Chemical reactions

Chemistry Chapter C6: Chemical Reactions (and C5: Chemical Energetics)

Hello future chemist! Welcome to the exciting chapter on Chemical Reactions. This topic is the core of chemistry—it’s all about how substances change, why they change, and how quickly those changes happen. Everything from cooking an egg to generating electricity relies on chemical reactions, so mastering this chapter is key to understanding the world around you! Let's dive in!

Section 1: Physical vs. Chemical Changes (C6.1)

Before we explore complex reactions, we must distinguish between two fundamental types of changes matter undergoes.

What is a Physical Change?

A physical change alters the appearance or form of a substance but does not create a new chemical substance. The chemical structure remains the same.

• Example: Melting ice. Ice (solid \(H_2O\)) turns into water (liquid \(H_2O\)). It’s still water, just in a different state.

Key features of Physical Changes:

• Usually easily reversible.
• No new chemical bonds are broken or formed.
• Includes changes of state (melting, freezing, boiling, condensing).

What is a Chemical Change?

A chemical change (or chemical reaction) results in the formation of one or more new substances with different chemical properties.

• Example: Burning wood. Wood reacts with oxygen to form ash, carbon dioxide, and water vapour. You cannot turn the ash back into wood simply by cooling it down.

Signs that a Chemical Reaction has occurred:

1. A permanent colour change (e.g., mixing two colourless liquids produces a coloured solid).
2. A gas is produced (often seen as effervescence or bubbling).
3. A solid (precipitate) is formed when two solutions are mixed.
4. A significant temperature change (heat given out or taken in).
5. Light or sound energy is produced.

Section 2: Chemical Energetics (C5.1)

All chemical reactions involve energy changes. They either release energy or absorb energy. This is called Chemical Energetics.

1. Exothermic Reactions

An exothermic reaction is a reaction that transfers thermal energy to the surroundings.

• This causes the temperature of the surroundings to increase. The container feels hot.
• Analogy: If you are running (a type of reaction), you release energy (heat) to the environment.
• Examples: Combustion (burning), neutralisation reactions, respiration.

2. Endothermic Reactions

An endothermic reaction is a reaction that takes in thermal energy from the surroundings.

• This causes the temperature of the surroundings to decrease. The container feels cold.
• Analogy: When you sweat, the evaporation process takes heat from your skin, cooling you down.
• Examples: Thermal decomposition, photosynthesis, dissolving certain salts like ammonium nitrate in water (used in instant cold packs).

Energy Profiles and Activation Energy (Supplement)

For a reaction to start, particles must collide with enough energy. This minimum required energy is called the activation energy (\(E_a\)).

The overall energy change in a reaction is the difference in energy between the reactants and the products.

Bond Breaking and Bond Making

• Bond breaking requires energy input (it is endothermic).
• Bond making releases energy (it is exothermic).

The overall energy change depends on whether more energy is released by making bonds than is absorbed by breaking them.

Exothermic Reaction: Energy released (bond making) > Energy absorbed (bond breaking). Overall energy transfer is out.

Endothermic Reaction: Energy absorbed (bond breaking) > Energy released (bond making). Overall energy transfer is in.

Interpreting Reaction Pathway Diagrams (Supplement)

Reaction pathway diagrams visually represent the energy changes during a reaction.

The diagram shows:

• The initial energy of the reactants.
• The final energy of the products.
• The peak of the curve represents the energy barrier, which is overcome by the activation energy (\(E_a\)).
• The difference between the reactants' energy and the products' energy is the overall energy change of reaction.

(Note: Since I cannot draw diagrams, imagine a simple graph with Energy on the Y-axis and Reaction Pathway on the X-axis.)

For Exothermic reactions: The energy level of the products is lower than the reactants (the overall energy change is negative, energy is lost to surroundings).

For Endothermic reactions: The energy level of the products is higher than the reactants (the overall energy change is positive, energy is taken from surroundings).

Section 3: Rate of Reaction (C6.2)

The rate of reaction measures how quickly reactants are used up or how quickly products are formed. A faster rate means the reaction finishes sooner.

We investigate the rate by monitoring changes over time, usually:

1. Measuring the volume of gas produced.
2. Measuring the change in mass of the reaction mixture (if a gas is escaping).
3. Measuring how quickly a solution turns cloudy (the 'disappearing cross' experiment).

Don't worry if measuring the rate seems tricky—it's just about watching how fast something changes!

Factors Affecting the Rate of Reaction (Core)

The rate of a chemical reaction can be changed by adjusting four key factors:

1. Concentration of Solutions

Increasing the concentration of reactants increases the rate of reaction.

2. Surface Area of Solids

Increasing the surface area of a solid reactant (e.g., crushing lumps into a powder) increases the rate of reaction.

• Analogy: A sugar cube dissolves slowly, but powdered sugar dissolves quickly because more particles are exposed to the water immediately.

3. Temperature

Increasing the temperature increases the rate of reaction.

4. Catalyst

Adding a catalyst increases the rate of reaction.

• A catalyst is a substance that speeds up a chemical reaction but remains chemically unchanged at the end of the reaction.
• Catalysts work by providing an alternative reaction pathway that has a lower activation energy (\(E_a\)).
• Did you know? Enzymes (B5) are biological catalysts!

Section 4: Explaining Rate using Collision Theory (Supplement)

To understand why these factors change the rate, we use Collision Theory.

Collision theory states that particles (atoms, molecules, or ions) must collide with each other with the correct orientation and sufficient energy (greater than or equal to the activation energy, \(E_a\)) to react. These successful collisions are called effective collisions.

How Factors Affect Effective Collisions (Supplement)

The rate of reaction is proportional to the frequency of effective collisions.

A. Effect of Concentration (Number of particles per unit volume)

• When concentration increases, there are more particles per unit volume.
• This means particles are packed closer together, leading to a higher frequency of collisions overall.
• Therefore, the frequency of effective collisions increases, and the rate speeds up.

B. Effect of Surface Area

• When a solid is broken up, the surface area exposed to the other reactant increases.
• This increases the number of sites where collisions can occur.
• This leads to a higher frequency of collisions, and the rate speeds up.

C. Effect of Temperature (Kinetic Energy of Particles)

• When temperature increases, the particles gain kinetic energy and move faster.
• This results in two effects:
  1. The particles collide more frequently (higher frequency of collisions).
  2. A much larger proportion of collisions have energy greater than the activation energy (\(E_a\)).
• The second point is the most important reason for the increased rate.

D. Effect of a Catalyst (Activation Energy, \(E_a\))

• A catalyst works by providing an alternative reaction pathway with a lower activation energy, \(E_a\).
• Since less energy is required for a collision to be effective, more of the particles already have enough energy to react.
• This greatly increases the frequency of effective collisions, speeding up the rate. The catalyst itself is not used up.

Section 5: Redox Reactions (C6.3)

Many chemical reactions involve the transfer of oxygen. These are known as Redox Reactions. The term 'Redox' is a combination of Reduction and Oxidation.

Definitions based on Oxygen Transfer (Core)

1. Oxidation

Oxidation is defined as the gain of oxygen by a substance.

• Example: Magnesium metal reacting with oxygen to form magnesium oxide: $$2Mg + O_2 \rightarrow 2MgO$$ Magnesium has gained oxygen, so it has been oxidised.

2. Reduction

Reduction is defined as the loss of oxygen by a substance.

• Example: Copper(II) oxide reacting with hydrogen: $$CuO + H_2 \rightarrow Cu + H_2O$$ Copper(II) oxide has lost oxygen, so it has been reduced.

Simultaneous Oxidation and Reduction

In every redox reaction, oxidation and reduction happen simultaneously (at the same time). You can’t have one without the other!

Let's look at the reaction between copper(II) oxide and hydrogen again: $$CuO + H_2 \rightarrow Cu + H_2O$$

• \(CuO\) loses oxygen to become \(Cu\). \(CuO\) is reduced.
• \(H_2\) gains oxygen to become \(H_2O\). \(H_2\) is oxidised.

Important Naming Convention (Core)

You need to identify oxidation and reduction based on the names of ions. This is important when dealing with transition metals like iron and copper, which can exist in multiple ion states:

• Iron(II) ion (\(Fe^{2+}\)) is oxidised to Iron(III) ion (\(Fe^{3+}\)) (it loses an electron).
• Copper(II) ion (\(Cu^{2+}\)) is reduced to Copper(I) ion (\(Cu^{+}\)) or to Copper metal (it gains electrons).

In the Combined Science syllabus, you must be able to recognise these different states, even if the definition provided in C6.3 focuses primarily on oxygen transfer.