Chemistry Study Notes: Equilibrium
Hey everyone! Welcome to your study notes on Chemical Equilibrium. Don't worry if this topic sounds a bit intimidating at first. We're going to break it down together. Think of it like a perfectly balanced tug-of-war. We'll explore what happens when chemical reactions don't just go in one direction, how we can describe this balance with a special number called Kc, and how reactions respond when we try to change the conditions using Le Chatelier’s Principle. Understanding this is super important, especially for making useful chemicals in industry!
1. Dynamic Equilibrium: The Balancing Act
What are Reversible Reactions?
Most reactions you've learned about so far go in one direction. For example, when you burn magnesium, it turns into magnesium oxide. It doesn't easily turn back. These are irreversible reactions.
However, many reactions can go both forwards and backwards at the same time! These are called reversible reactions. We show this using a special double arrow symbol (⇌) instead of a single arrow (→).
Example: The reaction between nitrogen and hydrogen to form ammonia.
$$N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$$
This means that while nitrogen and hydrogen are reacting to form ammonia (the forward reaction), ammonia is also breaking down to form nitrogen and hydrogen (the reverse reaction).
What is Dynamic Equilibrium?
Imagine a busy shop. People are constantly entering, and other people are constantly leaving. If the rate of people entering is the same as the rate of people leaving, the total number of people inside the shop stays constant. It looks like nothing is changing from the outside, but there's a lot of activity happening!
This is exactly what dynamic equilibrium is. It's a state where:
- The rate of the forward reaction is equal to the rate of the reverse reaction.
- The concentrations of all the reactants and products remain constant.
- It's "dynamic" because both reactions are still happening, they're just perfectly balanced.
Important Conditions for Equilibrium:
- The system must be closed. This means nothing can get in or out (like a lid on a bottle).
- The macroscopic properties (like colour, concentration, pressure) do not change over time.
Key Takeaway
Dynamic equilibrium is a balanced state in a reversible reaction where the forward and reverse reactions occur at the same rate, resulting in constant concentrations of reactants and products. It’s a state of balance, not a state of stoppage!
2. The Equilibrium Constant (Kc): A Number That Tells a Story
So, we know that at equilibrium, the concentrations are constant. But are there more products or more reactants? The equilibrium constant (Kc) is a value that gives us the answer!
How to Write the Kc Expression
For any general reversible reaction at equilibrium:
$$aA + bB \rightleftharpoons cC + dD$$
The equilibrium constant expression is written as:
$$K_c = \frac{[C]^c[D]^d}{[A]^a[B]^b}$$
Let's break that down:
- The concentrations of the products are on the top (numerator).
- The concentrations of the reactants are on the bottom (denominator).
- The square brackets [ ] mean "concentration of" in mol dm⁻³.
- The stoichiometric coefficients (the numbers a, b, c, d from the balanced equation) become the powers for each concentration.
Step-by-Step Example: Haber Process
Let's write the Kc expression for the formation of ammonia:
$$N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$$
- Identify products: NH₃. Its coefficient is 2. So, the top is [NH₃]².
- Identify reactants: N₂ and H₂. Their coefficients are 1 and 3. So, the bottom is [N₂]¹[H₂]³.
- Put it together:
$$K_c = \frac{[NH_3]^2}{[N_2][H_2]^3}$$
Watch Out! A Common Mistake
It’s very easy to forget to use the coefficients as powers! Always double-check your balanced equation when writing a Kc expression. Also, remember it's always [Products] over [Reactants].
What Does the Value of Kc Mean?
The size of Kc tells you about the position of the equilibrium – whether it favours the products or the reactants.
- If Kc is large (Kc > 1): The top number ([Products]) is much bigger than the bottom number ([Reactants]). This means at equilibrium, there are more products than reactants. The equilibrium position lies to the right.
- If Kc is small (Kc < 1): The bottom number ([Reactants]) is much bigger. This means at equilibrium, there are more reactants than products. The equilibrium position lies to the left.
Crucial Point: The value of Kc for a given reaction is only constant at a constant temperature. If you change the temperature, you change the value of Kc!
Calculating Kc
To calculate Kc, you just need to know the concentrations of all reactants and products *at equilibrium* and plug them into the expression.
Example Calculation:
For the reaction $$H_2(g) + I_2(g) \rightleftharpoons 2HI(g)$$, the equilibrium concentrations in a sealed flask at 400°C were found to be:
[H₂] = 0.15 mol dm⁻³
[I₂] = 0.15 mol dm⁻³
[HI] = 1.70 mol dm⁻³
Calculate Kc.
Step 1: Write the Kc expression.
$$K_c = \frac{[HI]^2}{[H_2][I_2]}$$
Step 2: Substitute the equilibrium concentrations.
$$K_c = \frac{(1.70)^2}{(0.15)(0.15)}$$
Step 3: Calculate the value.
$$K_c = \frac{2.89}{0.0225} = 128.4$$
(Note: Kc usually has no units) Since Kc is much larger than 1, this tells us that at this temperature, the equilibrium strongly favours the formation of HI.
Key Takeaway
Kc is the ratio of [products] to [reactants] at equilibrium. A large Kc means the reaction favours products, while a small Kc means it favours reactants. Kc is constant unless the temperature changes.
3. Le Chatelier's Principle: Responding to Change
The Core Idea
What happens if we have a system at equilibrium and we suddenly disturb it (e.g., by adding more reactant)? The French chemist Henri Le Chatelier figured this out.
Le Chatelier’s Principle states that: "When a change is applied to a system at equilibrium, the system will shift in a way that counteracts the change."
Analogy: The Seesaw
Think of equilibrium as a perfectly balanced seesaw. If you add weight (a "change") to one side, it becomes unbalanced. To counteract this, the seesaw needs to shift some of that new weight to the other side to try and get balanced again. A chemical system does the same thing!
What Does "Shifting Equilibrium" Mean?
- A shift to the right means the forward reaction rate increases temporarily to produce more products.
- A shift to the left means the reverse reaction rate increases temporarily to produce more reactants.
Applying Le Chatelier's Principle
Let's use a visual example to see this principle in action. This reaction involves a colour change from pink to blue:
$$Co(H_2O)_6^{2+}(aq) + 4Cl^-(aq) \rightleftharpoons CoCl_4^{2-}(aq) + 6H_2O(l)$$
(Pink) (Blue)
1. Effect of Changing Concentration
The Rule: The system tries to get rid of what you add, and replace what you remove.
- Change: Add more Cl⁻ ions (e.g., by adding some concentrated HCl).
- System's Response: "Whoa, too much Cl⁻! I need to use it up."
- How: It will favour the forward reaction to consume the extra Cl⁻.
- Result: The equilibrium shifts to the right. The solution will turn more blue.
- Change: Remove CoCl₄²⁻ (this is hard to do, but imagine we could).
- System's Response: "Hey, where did the CoCl₄²⁻ go? I need to make more."
- How: It will favour the forward reaction to replace the lost CoCl₄²⁻.
- Result: The equilibrium shifts to the right.
Important Point: Changing the concentration causes the equilibrium position to shift, but the value of Kc does NOT change (as long as the temperature stays the same).
2. Effect of Changing Temperature
The Rule: Think of 'heat' as a reactant (endothermic) or a product (exothermic).
Let's say our example reaction is endothermic, meaning it absorbs heat. (ΔH is positive).
$$Heat + Co(H_2O)_6^{2+}(aq) + 4Cl^-(aq) \rightleftharpoons CoCl_4^{2-}(aq) + 6H_2O(l)$$
- Change: Increase the temperature (add heat).
- System's Response: "It's getting hot in here! I need to use up this extra heat."
- How: It will favour the endothermic reaction (the one that consumes heat).
- Result: The equilibrium shifts to the right. The solution will turn more blue.
- Change: Decrease the temperature (remove heat by putting it in an ice bath).
- System's Response: "Brrr, it's cold! I need to produce some heat."
- How: It will favour the exothermic reaction (the one that releases heat).
- Result: The equilibrium shifts to the left. The solution will turn more pink.
Important Point: Changing the temperature is the ONLY thing that changes the value of Kc.
- For an endothermic reaction, increasing T shifts it right, making more products. So, Kc increases.
- For an exothermic reaction, increasing T shifts it left, making fewer products. So, Kc decreases.
Did you know?
The Haber Process ($$N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$$) is exothermic. According to Le Chatelier's principle, a low temperature would give the highest yield of ammonia. However, at low temperatures, the reaction is incredibly slow! So, industries use a compromise temperature (around 450°C) and a catalyst to get a reasonable yield at a reasonable rate. It's a real-life balancing act!
Quick Review: Le Chatelier's Principle Summary
- Change Concentration: System shifts to use up what's added or replace what's removed. Kc is unchanged.
- Change Temperature: System shifts to favour the endothermic direction if heated, or the exothermic direction if cooled. Kc changes.
You've Got This! Final Summary
That's the core of chemical equilibrium! It might take a little practice, but you can master it. Just remember these key ideas:
- Dynamic Equilibrium: Not static! Forward rate = Reverse rate. Concentrations are constant.
- Kc (Equilibrium Constant): A ratio of [Products] / [Reactants]. Its value tells you which side is favoured at equilibrium.
- Le Chatelier's Principle: The system counteracts any change you make to it. Use this principle to predict how equilibrium will shift.
Keep practising writing Kc expressions and applying Le Chatelier's principle to different scenarios. You can do it!