👋 Welcome to Physical Chemistry: How Reactions Happen!
Hey there! Chemistry isn't just about what substances are; it's also about how fast they change and where the energy goes when they react. This chapter—Physical Chemistry—is about understanding the rules of these chemical changes.
Don't worry if words like "equilibrium" sound intimidating! We'll break down these concepts using simple language and everyday examples. By the end of this section, you'll know exactly how to speed up a reaction and why some reactions feel hot while others feel cold. Let's dive in!
1. Rates of Reaction
1.1. What is the Rate of Reaction?
The rate of reaction is simply how quickly reactants are used up and products are formed. A fast reaction (like an explosion) has a high rate; a slow reaction (like rusting) has a low rate.
The Collision Theory (The Golden Rule)
For any reaction to occur, the particles (atoms, molecules, or ions) of the reactants must collide. But not just any bump will do! Two things must be true for a collision to be successful:
- They must collide in the correct orientation (they must "hit" the right spot).
- They must collide with enough energy. This minimum energy required is called the Activation Energy (Ea).
Analogy: Imagine trying to light a match. You need the match head (reactant A) to hit the striking surface (reactant B) at the right angle, AND you need to rub it hard enough (enough activation energy) to start the fire.
1.2. Factors Affecting the Rate
We can control the rate of a reaction by changing conditions that increase the frequency of successful collisions. There are four main factors you need to know:
Quick Mnemonic: C-A-T-S
A. Concentration (for solutions) or Pressure (for gases)
If you increase the concentration of a reactant, you pack more particles into the same volume.
- Effect: There are more particles, so collisions happen more often.
- Result: Increased frequency of successful collisions, therefore a faster rate.
Think of a busy market versus an empty street. You are much more likely to bump into someone in the busy market (high concentration).
B. Surface Area (for solids)
If a reactant is a solid, only the particles on the outside surface can react.
- Increase Surface Area: Crush the solid into a powder.
- Effect: More reactant particles are exposed to the other reactants.
- Result: More collision sites are available, leading to a faster rate.
Example: Powdered sugar dissolves much faster than a sugar cube because it has a massive surface area exposed to the water.
C. Temperature
Increasing the temperature affects the reaction rate in two important ways:
- Frequency: Particles move faster, so they collide more frequently.
- Energy: A larger proportion of particles now have energy equal to or greater than the Activation Energy (Ea).
This second point is the most important reason why increasing temperature speeds up a reaction.
D. Catalysts
A catalyst is a substance that speeds up a chemical reaction but is chemically unchanged itself at the end of the reaction.
- How they work: Catalysts provide an alternative reaction pathway that has a lower Activation Energy (Ea).
- Benefit: Because the energy needed is lower, many more of the normal collisions now become successful collisions, speeding up the rate dramatically.
Did you know? Catalytic converters in cars use precious metals (catalysts) to turn harmful exhaust gases into less harmful ones quickly.
1.3. Measuring the Rate of Reaction
To measure rate, we must observe how quickly a reactant changes or a product forms. Common methods include:
- Measuring Volume of Gas Produced: Collect the gas in a syringe or inverted measuring cylinder and record the volume over time.
- Measuring Mass Loss: If a gas is produced, the system loses mass. Placing the reaction vessel on a balance and recording the mass drop over time measures the rate.
- Timing how long it takes for a visual change to occur (e.g., disappearance of a precipitate, colour change).
🔑 Quick Review: Rates of Reaction
Rate = Speed of reaction. Must have successful collisions (enough energy, right orientation).
4 Factors: Concentration, Surface Area, Temperature, Catalyst.
Catalysts: Speed up rate by lowering Activation Energy.
2. Energetics: Heat Changes in Reactions
Every chemical reaction involves an energy change, usually in the form of heat. This energy change happens because energy is needed to break bonds in the reactants and energy is released when new bonds are made in the products.
2.1. Exothermic Reactions
In an exothermic reaction, the total energy released when making new bonds is greater than the energy required to break the old bonds.
- Effect: Energy is released (usually as heat) into the surroundings.
- Observation: The temperature of the surroundings increases (it feels hot).
- Mnemonic: Exo means Exit. Heat exits the system.
Examples: Combustion (burning fuel), neutralisation reactions, respiration (in living things).
2.2. Endothermic Reactions
In an endothermic reaction, the energy required to break the original bonds is greater than the energy released when making new bonds.
- Effect: Energy is taken in (absorbed) from the surroundings.
- Observation: The temperature of the surroundings decreases (it feels cold).
- Mnemonic: Endo means Enter. Heat enters the system (from the surroundings).
Examples: Thermal decomposition reactions, dissolving certain salts in water (used in chemical cold packs).
2.3. Energy Profiles
We can visualise the energy changes using simple diagrams:
Energy Profile Diagram (General Features):
- The starting point shows the energy of the reactants.
- The peak of the curve represents the energy state needed to start the reaction (the energy barrier). The height from the reactants to the peak is the Activation Energy (Ea).
- The final point shows the energy of the products.
For an Exothermic Reaction: The energy level of the products is lower than the energy level of the reactants. Energy has been lost to the surroundings.
For an Endothermic Reaction: The energy level of the products is higher than the energy level of the reactants. Energy has been absorbed from the surroundings.
(Note: You should be able to sketch and label these two types of diagrams.)
3. Reversible Reactions and Equilibrium
Most reactions we study go to completion (all reactants turn into products). However, some reactions are capable of going both ways. These are called reversible reactions.
3.1. Reversible Reactions
A reversible reaction is one where the products can react together to reform the original reactants.
- We show this using a special double arrow symbol: \( \rightleftharpoons \)
Example: The Haber process (making ammonia) is reversible:
\( \text{N}_2 \text{(g)} + 3\text{H}_2 \text{(g)} \rightleftharpoons 2\text{NH}_3 \text{(g)} \)
The reaction going left-to-right is the forward reaction. The reaction going right-to-left is the reverse reaction.
3.2. Chemical Equilibrium
If a reversible reaction takes place in a closed system (meaning nothing can escape or be added), it will eventually reach a state called chemical equilibrium.
What happens at Equilibrium?
Equilibrium does not mean the reaction has stopped! It means the speed of the forward reaction is exactly equal to the speed of the reverse reaction.
Rate of Forward Reaction = Rate of Reverse Reaction
- At equilibrium, the concentration of reactants and products stays constant.
- It looks like nothing is happening, but both reactions are still occurring constantly—this is known as dynamic equilibrium.
Analogy: Imagine two escalators side-by-side, one going up and one going down, connecting Floor 1 and Floor 2. If 10 people go up every minute, and 10 people come down every minute, the number of people on Floor 1 and Floor 2 stays constant. That is dynamic equilibrium!
🚀 Final Boost!
Physical Chemistry ties the whole subject together by explaining the why and how fast. Remember the key concepts: collisions must be successful (Rates), energy must be conserved (Energetics), and some reactions are never finished (Equilibrium).
Keep practising those definitions and you will master this section!