Chemistry Study Notes: Effects of Temperature and Concentration

Hello there! Welcome to your study notes on a super important topic in Chemistry. Ever wondered why food cooks faster at a higher heat, or why a concentrated cleaning liquid works better than a diluted one? It all comes down to the effects of temperature and concentration on chemical reactions.

In these notes, we'll explore two main ideas:

  1. How these factors change the SPEED (rate) of a reaction.
  2. How they can shift the BALANCE (equilibrium) of a reversible reaction.

Don't worry if these terms seem tricky at first. We'll break everything down with simple language, everyday examples, and helpful analogies. Let's get started!


Part 1: Changing the SPEED of a Reaction (Rate of Reaction)

What is Reaction Rate and Why Do Collisions Matter?

Before we dive in, let's quickly review the basics. The rate of reaction is just a fancy way of saying how fast reactants are turned into products.

For a reaction to happen, reactant particles (atoms, ions, or molecules) must do two things:

  1. Collide with each other.
  2. Collide with enough energy to break old bonds and form new ones. This minimum energy needed is called the activation energy (Ea).

This is known as the Collision Theory. Think of it as the fundamental rulebook for reaction speed. Any factor that increases the number of successful (or 'effective') collisions per second will speed up the reaction.

Quick Review Box

Reaction Rate: The speed of a reaction.
Effective Collision: A collision that has enough energy (≥ activation energy) and the correct orientation to result in a reaction.
To speed up a reaction: Increase the frequency of effective collisions!


Factor 1: The Effect of Concentration

What Happens?

Increasing the concentration of reactants increases the rate of reaction.

Why Does It Happen?

Concentration refers to how many reactant particles are packed into a certain volume. A higher concentration means more particles are squeezed into the same space.

  • When particles are more crowded, they are going to bump into each other much more often.
  • This leads to more frequent collisions between reactant particles.
  • With more collisions happening every second, the number of effective collisions per second also increases.
Analogy: The Crowded MTR

Imagine you're in an MTR station. During off-peak hours (low concentration), you can walk around without bumping into anyone. But during rush hour (high concentration), the platform is packed, and you can barely move without colliding with other people. It's the same for chemical particles!

Real-World Example

A concentrated acid reacts much more violently and quickly with a metal than a dilute acid. This is because there are more acid particles in the same volume of water to collide with the metal surface.

Key Takeaway for Concentration

Higher Concentration → More particles in the same volume → More frequent collisions → Faster Reaction Rate.


Factor 2: The Effect of Temperature

What Happens?

Increasing the temperature significantly increases the rate of reaction.

A common rule of thumb is that for many reactions, a 10°C rise in temperature roughly doubles the reaction rate!

Why Does It Happen? (This has TWO reasons!)

This is a really important concept, and students often only remember the first reason. Make sure you understand both!

Reason 1: Particles move faster.

  • When you heat a substance, you give its particles more kinetic energy.
  • This makes them move around much faster.
  • Faster-moving particles will collide more often (more frequent collisions).

Reason 2: Collisions are more energetic. (THE MAIN REASON!)

  • This is the bigger and more important effect.
  • At a higher temperature, a much larger proportion of the particles have energy equal to or greater than the activation energy (Ea).
  • This means that the percentage of collisions that are successful (effective) dramatically increases.

So, not only are there more collisions, but a much larger fraction of those collisions actually lead to a reaction!

Analogy: Throwing Balls Over a Wall

Imagine the activation energy is a high wall. You and your friends (the particles) are trying to throw balls (energy) over it.
Low Temperature: Everyone is tired and throwing gently. The balls hit the wall but don't go over. Very few successful "reactions".
High Temperature: Everyone has a lot of energy and is throwing the balls hard. A much higher percentage of the balls now have enough energy to clear the wall. The rate of "success" shoots up!

Key Takeaway for Temperature

Higher Temperature → Particles move faster AND have more energy → More frequent collisions AND a higher percentage of collisions are effective → Much Faster Reaction Rate.



Part 2: Shifting the BALANCE of a Reaction (Chemical Equilibrium)

What is Chemical Equilibrium?

Some reactions don't just go one way. They are reversible, meaning the products can react to form the reactants again. We show this with a special double arrow:

$$ \text{Reactants} \rightleftharpoons \text{Products} $$

Dynamic Equilibrium is reached when:

  • The rate of the forward reaction (Reactants → Products) is equal to the rate of the backward reaction (Products → Reactants).
  • The concentrations of all reactants and products remain constant.

Important: The reactions have NOT stopped! They are still happening, but they are perfectly balanced.

Analogy: The Escalator

Imagine two people on an escalator. One is walking up at exactly the same speed the other is walking down. From a distance, it looks like their positions aren't changing (constant concentration), but they are both in constant motion (dynamic). That's equilibrium!


Le Chatelier's Principle: The Golden Rule

When a system at equilibrium is disturbed by a change (like changing temperature or concentration), it will behave in a way that helps to partially counteract the change. It always tries to restore the balance.

Common Mistake to Avoid

Students often forget the word "partially". The system can't completely undo the change, but it does its best to oppose it.


The Effect of Concentration on Equilibrium

The Rule

If you change the concentration of one substance in an equilibrium mixture, the system will shift to counteract that change.

  • If you ADD more of a substance, the equilibrium will shift to the other side to use it up.
  • If you REMOVE a substance, the equilibrium will shift towards that substance's side to make more of it.
Step-by-Step Example

Let's look at the famous equilibrium between pink cobalt ions and blue cobalt ions:

$$ \underbrace{Co(H_2O)_6^{2+}(aq)}_{\text{Pink}} + 4Cl^-(aq) \rightleftharpoons \underbrace{CoCl_4^{2-}(aq)}_{\text{Blue}} + 6H_2O(l) $$

Scenario 1: We add more chloride ions (Cl⁻).

  1. The Change: Concentration of a reactant (Cl⁻) has increased.
  2. The Counteraction: The system wants to decrease the concentration of Cl⁻.
  3. The Shift: To use up the extra Cl⁻, the forward reaction must speed up. The equilibrium position shifts to the right.
  4. The Observation: The solution will become more blue.

Scenario 2: We add water (H₂O), which dilutes everything and effectively removes some CoCl₄²⁻.

  1. The Change: Concentration of a product (CoCl₄²⁻) has decreased.
  2. The Counteraction: The system wants to make more CoCl₄²⁻.
  3. The Shift: To make more CoCl₄²⁻, the forward reaction must be favoured. But wait, adding water is a bit more complex. The easiest way to think about it for this specific reaction is that you added a product (H₂O). The system wants to use up the extra H₂O, so it shifts to the left.
  4. The Observation: The solution will become more pink.
Did you know?

Changing the concentration will shift the position of the equilibrium, but it has NO EFFECT on the value of the equilibrium constant (Kc) as long as the temperature stays the same.


The Effect of Temperature on Equilibrium

The Rule

To figure this out, you need to know if the reaction is exothermic (releases heat, ΔH is negative) or endothermic (absorbs heat, ΔH is positive). The easiest way is to treat 'heat' as a reactant or a product!

  • If you INCREASE the temperature (add heat), the equilibrium will shift in the endothermic direction to absorb the extra heat.
  • If you DECREASE the temperature (remove heat), the equilibrium will shift in the exothermic direction to produce more heat.
Step-by-Step Example

Let's consider the production of ammonia in the Haber Process. The forward reaction is exothermic.

$$ N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g) \quad \Delta H = -92 \, \text{kJ/mol} $$

Since it's exothermic, we can write 'heat' as a product:

$$ N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g) + \text{HEAT} $$

Scenario 1: We increase the temperature.

  1. The Change: We are adding heat.
  2. The Counteraction: The system wants to get rid of the extra heat.
  3. The Shift: To use up heat, the reaction must go in the reverse (endothermic) direction. The equilibrium position shifts to the left.
  4. The Result: The yield of ammonia (NH₃) will decrease.

Scenario 2: We decrease the temperature.

  1. The Change: We are removing heat.
  2. The Counteraction: The system wants to produce more heat to warm itself up.
  3. The Shift: To produce heat, the reaction must go in the forward (exothermic) direction. The equilibrium position shifts to the right.
  4. The Result: The yield of ammonia (NH₃) will increase.
Very Important Point!

Temperature is the ONLY factor that changes the value of the equilibrium constant (Kc).
- For an exothermic forward reaction, increasing temperature decreases Kc.
- For an endothermic forward reaction, increasing temperature increases Kc.


Final Summary Table
| Change Made | Effect on Reaction RATE | Effect on Equilibrium POSITION | | :--- | :--- | :--- | | Increase Concentration of a Reactant | Increases Rate | Shifts to the right (towards products) | | Decrease Concentration of a Reactant | Decreases Rate | Shifts to the left (towards reactants) | | Increase Temperature | Greatly Increases Rate | Shifts in the endothermic direction | | Decrease Temperature | Greatly Decreases Rate | Shifts in the exothermic direction |

Great job making it through these notes! Keep reviewing the analogies and examples. Understanding how we can control chemical reactions by tweaking these conditions is a core skill in chemistry and helps explain so much of the world around us. Keep up the great work!