Understanding Chemical Equilibrium – A Guide for Grade 12 Physical Sciences Learners

Dear Grade 12 learner

If you’ve ever read about chemical equilibrium and felt overwhelmed by the formulas, arrows, or long definitions — you're not alone.

This topic can feel like a balancing act.

But once you understand what’s really going on, it all starts to make sense.

Let's walk through it together — gently, step by step — so you can face it with confidence.

🧪 What is Chemical Equilibrium?

In a reversible reaction, the products can react to form the reactants again.

Over time, the forward and reverse reactions happen at the same rate — and that’s when the system reaches chemical equilibrium.

Chemical equilibrium is when the rate of the forward reaction equals the rate of the reverse reaction, and the concentrations of reactants and products remain constant (but not necessarily equal).

🔁 A Simple Example

Let’s say we have the following reversible reaction:

N2(g)+3H2(g)⇌2NH3(g)\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons

2\text{NH}_3(g)N2​(g)+3H2​(g)⇌2NH3​(g)

This is the Haber Process — used to make ammonia for fertilizers.

At first, nitrogen and hydrogen react to make ammonia (NH₃).

But after a while, some of the ammonia molecules start breaking down to form nitrogen and hydrogen again.

Eventually, the forward reaction (making NH₃) and the reverse reaction (breaking NH₃) happen at the same rate. That’s equilibrium!

⚖️ Important Characteristics of Equilibrium

1. It’s dynamic – reactions are still happening, just at equal rates.

2. Concentrations stay constant – but they’re not necessarily the same for all substances.

3. Only occurs in closed systems – where no substances can escape.

4. Can be disturbed – and that’s where Le Chatelier’s Principle comes in.

   
   

🧠 Le Chatelier’s Principle (Simplified)

This principle says:

If something is changed in a system at equilibrium, the system will shift to oppose the change and try to restore balance.

📘 Example 1: Changing Concentration

N2+3H2⇌2NH3\text{N}_2 + 3\text{H}_2 \rightleftharpoons 2\text{NH}_3N2​+3H2​⇌2NH3​

What if we add more N₂?

The system wants to use it up, so it shifts to the right — making more NH₃.

What if we remove NH₃?

The system shifts to the right to replace it.

🌡️ Example 2: Changing Temperature

If the forward reaction is exothermic (releases heat), increasing the temperature will shift the equilibrium to the left — favouring the reverse (endothermic) reaction.

Tip: Think of heat like a reactant in endothermic reactions and a product in exothermic ones.

🧮 The Equilibrium Constant (Kc)

Kc tells us how much product or reactant is present at equilibrium.

The general formula is:

Kc=[Products][Reactants]K_c

= \frac{[\text{Products}]}{[\text{Reactants}]}Kc​=[Reactants][Products]​

Let’s use a simple example:

H2(g)+I2(g)⇌2HI(g)\text{H}_2(g) + \text{I}_2(g) \rightleftharpoons

2\text{HI}(g)H2​(g)+I2​(g)⇌2HI(g)Kc=[HI]2[H2][I2]K_c

= \frac{[\text{HI}]^2}

{[\text{H}_2][\text{I}_2]}Kc​=[H2​][I2​][HI]2​

Kc is a fixed value at a specific temperature.

• If Kc > 1 → more products

• If Kc < 1 → more reactants

You're not expected to memorize Kc values — just know how to calculate them and what they mean.

🧠 Quick Recap:

Concept	Meaning
Reversible Reaction	A reaction that goes forward and backward
Chemical Equilibrium	Forward and reverse reactions happen at the same rate
Le Chatelier’s Principle	The system shifts to oppose any change
Kc	A constant showing the ratio of products to reactants at equilibrium

💡 A Tip from One Learner to Another

Try not to just memorise the theory.

Think of equilibrium as a conversation between reactants and products — always adjusting to keep the peace.

Understanding why the system shifts makes the math and theory easier to follow.

📖 Practice Question

Let’s say at equilibrium:

• [H2]=0.4 mol/L[\text{H}_2] = 0.4 \text{ mol/L}[H2​]=0.4 mol/L

• [I2]=0.4 mol/L[\text{I}_2] = 0.4 \text{ mol/L}[I2​]=0.4 mol/L

• [HI]=0.8 mol/L[\text{HI}] = 0.8 \text{ mol/L}[HI]=0.8 mol/L

  
  

Kc=[HI]2[H2][I2]=(0.8)20.4×0.4=0.640.16=4.0K_c

= \frac{[\text{HI}]^2}{[\text{H}_2][\text{I}_2]} = \frac{(0.8)^2}{0.4

\times 0.4} = \frac{0.64}{0.16}

= 4.0Kc​=[H2​][I2​][HI]2​=0.4×0.4(0.8)2​=0.160.64​=4.0

So, Kc = 4.0, meaning there are more products than reactants at equilibrium.

🌟 Final Encouragement

Yes, chemical equilibrium can be challenging — but so are many things you’ve already overcome.

You're not just memorising chemical rules — you're learning how systems maintain balance, how nature adjusts itself, and how industries make everyday products.

Physical Science-chemical equilibrium.pdf

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Take your time.

Ask for help when needed.

Practice a little every day.

You can understand this — and you’re not alone on the journey.

You’ve got this, future scientist! 🧪⚖️💛