Rates of reaction and energy changes

Welcome to Rates of Reaction and Energy Changes!

In this chapter, we are going to explore two very important questions in Chemistry: How fast does a reaction go? and How much energy does it give out or take in? This is a key part of your Paper 4: Chemistry 2 exam. Whether you’re a science pro or find it a bit confusing, these notes will help you master the basics and the tricky bits. Let's dive in!

1. What is the "Rate" of Reaction?

The rate of reaction is simply a measure of how quickly reactants turn into products. Think of it like a race: how much distance is covered in a certain amount of time? In chemistry, it's how much product is made (or reactant used up) in a certain time.

Measuring the Rate

There are two main ways we measure this in the lab (as seen in your Core Practicals):
- Measuring gas production: If a reaction makes a gas (like reacting marble chips with hydrochloric acid), we can collect the gas in a syringe and see how much is made every 10 seconds.
- Observing a color change: In the "disappearing cross" experiment (sodium thiosulfate and hydrochloric acid), we time how long it takes for a solution to turn cloudy enough to hide a cross drawn under the flask.

Interpreting Reaction Graphs

When you look at a graph of Volume of Gas vs Time:
- Steep line: The reaction is very fast.
- Line getting less steep: The reaction is slowing down as reactants get used up.
- Flat line: The reaction has finished because one of the reactants has completely run out.

Quick Review:
Rate = Amount of reactant used or product formed / Time
The steeper the graph, the faster the reaction!

2. Collision Theory: The "How" of Reactions

Don't worry if this sounds fancy—it’s actually very simple! For a chemical reaction to happen, particles must collide with each other. But just bumping into each other isn't enough. They need two things:
1. To collide with enough energy (this is called the activation energy).
2. To collide frequently.

Analogy: Imagine "Bumper Cars." If the cars just tap each other gently, nothing happens. But if they crash together with a lot of speed (energy), they might cause a dent (a reaction)! To get more dents, you need more cars on the track (higher frequency) or faster cars (higher energy).

Key Takeaway: To increase the rate of reaction, you must increase the frequency of collisions or the energy of the collisions.

3. Factors Affecting the Rate

There are four main ways we can speed up a reaction. Here is how they work in terms of Collision Theory:

Temperature

When we increase the temperature, particles move faster. This means:
- They collide more frequently.
- They collide with more energy, so more collisions are "successful."

Concentration (and Pressure for gases)

Increasing concentration (more particles in a liquid) or pressure (squashing gas particles closer together) means there are more particles in the same space.
- This increases the frequency of collisions.

Surface Area

If you have a solid reactant (like a marble chip), breaking it into smaller pieces increases the surface area to volume ratio.
- More "inside" particles are now on the "outside" and available to react.
- This increases the frequency of collisions.

Example: Powdered sugar dissolves much faster in tea than a sugar cube because the powder has a huge surface area!

Quick Review Box:
- Higher Temp: Faster particles + more energetic collisions.
- Higher Conc/Pressure: More particles = more crowding = more collisions.
- More Surface Area: More exposed bits = more collisions.

4. Catalysts: The Shortcuts

A catalyst is a substance that speeds up a reaction without being used up itself. It is still there, unchanged, at the end of the reaction.

How do they work?

They provide an alternative "pathway" for the reaction that has a lower activation energy.
Analogy: If you want to get to the other side of a mountain, the "activation energy" is climbing over the top. A catalyst is like a tunnel through the mountain. It's a faster, easier way to get to the same destination!

Enzymes

Enzymes are biological catalysts. We use them in everyday life, such as using yeast to produce the alcohol in drinks. They work just like chemical catalysts but are made by living cells.

Key Takeaway: Catalysts save energy and time because they lower the "energy hill" (activation energy) that particles need to climb to react.

5. Energy Changes: Exothermic and Endothermic

During a chemical reaction, heat energy is usually moved around. We can measure this by looking at the temperature change of the surroundings (like the water in a flask).

Exothermic Reactions

Exothermic reactions give OUT heat to the surroundings.
- How to remember: Exothermic = Exit (heat leaves).
- The temperature of the surroundings increases (gets hotter).
- Examples: Combustion (burning), Neutralisation, and many Displacement reactions.

Endothermic Reactions

Endothermic reactions take IN heat from the surroundings.
- How to remember: Endothermic = Enter (heat enters the reaction).
- The temperature of the surroundings decreases (gets colder).
- Examples: Thermal decomposition, and the reaction between citric acid and sodium hydrogencarbonate.

Did you know? Sports ice packs use an endothermic reaction! When you squeeze the pack, chemicals mix and take in heat, making the pack feel freezing cold instantly.

6. The Science of Bonds (MEXO BENDO)

Every chemical reaction involves breaking bonds in the reactants and making new bonds in the products. This is where the energy change comes from.

1. Bond Breaking is Endothermic. It requires energy to pull atoms apart (Think: You have to put in effort to break a stick).
2. Bond Making is Exothermic. Energy is released when new bonds form.

Memory Aid: MEXO BENDO
- Making = Exothermic
- Breaking = Endothermic

Calculating the Energy Change

You might be asked to calculate the overall energy change using bond energy values given in a table. Use this simple "Left minus Right" rule:
Energy change = (Total energy needed to break bonds) – (Total energy released making bonds)

- If the answer is negative, the reaction is Exothermic (more energy was released than taken in).
- If the answer is positive, the reaction is Endothermic (more energy was taken in than released).

7. Reaction Profiles

A reaction profile is a graph that shows the energy levels of the reactants and products.

Exothermic Profile

- The Reactants are higher than the Products.
- This is because energy was given out to the surroundings.

Endothermic Profile

- The Reactants are lower than the Products.
- This is because energy was taken in from the surroundings.

Important Labels

On these graphs, you must be able to identify:
- Activation Energy: The "hump" on the graph. It is the energy measured from the reactants to the very top of the curve.
- Overall Energy Change: The difference in height between the reactants and the products.

Common Mistake to Avoid: When drawing activation energy, always start the arrow at the reactant line and go to the peak of the curve. Don't start from the bottom of the graph!

Final Key Takeaway:
- Exothermic: Products have LESS energy than reactants (Temperature goes UP).
- Endothermic: Products have MORE energy than reactants (Temperature goes DOWN).
- Activation Energy: The minimum energy needed for a collision to result in a reaction.