Effect of Concentration Change

The Core Idea: Disturbing Equilibrium by Changing Concentration

When a chemical system is at equilibrium, the forward and reverse reaction rates are equal, and the concentrations of reactants and products are constant. If you disturb this balance by adding or removing a reactant or product, the system will respond to counteract that change. This is a direct application of Le Chatelier’s principle.

The principle predicts two clear outcomes for concentration changes:

• Adding a substance: The system will shift in the direction that consumes the added substance.
• Removing a substance: The system will shift in the direction that replenishes the removed substance.

In other words, the composition of the equilibrium mixture changes to minimise the effect of the concentration change you made.

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How the Reaction Quotient Explains the Shift

The reaction quotient, $Q_c$, is the tool that lets us see why the shift happens. For the reaction:

The equilibrium constant is $K_c = \frac{[\text{HI}]^2}{[\text{H}_2][\text{I}_2]}$.

At equilibrium, $Q_c = K_c$.

Scenario: Adding $\text{H}_2$ to the equilibrium mixture.

1. The Disturbance: You suddenly increase $[\text{H}_2]$.
2. The Immediate Effect: The denominator of the $Q_c$ expression becomes larger. This makes the value of $Q_c$ smaller than $K_c$.

3. The System's Response: The system is now in a state where $Q_c < K_c$. To reach equilibrium again, the reaction must produce more products and consume reactants. This means the net reaction proceeds in the forward direction (towards products).
4. The New Equilibrium: The forward reaction consumes some of the added $\text{H}_2$ and some $\text{I}_2$ to produce more $\text{HI}$. The new equilibrium concentrations will be such that $Q_c = K_c$ again. The final concentration of $\text{H}_2$ will be higher than it was originally, but lower than it was immediately after the addition.

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A Powerful Commercial Application: Removing a Product

The principle of shifting equilibrium by removing a product is a cornerstone of industrial chemistry. If a reaction produces a gas or a volatile substance, continuously removing that substance from the reaction mixture keeps $Q_c$ perpetually less than $K_c$. This drives the reaction to completion, maximising the yield of the desired product.

Two classic examples from the textbook:

• Manufacture of Ammonia (Haber Process): In the reaction $\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g)$, ammonia is a gas. It is continuously liquefied and removed from the reaction mixture. This shifts the equilibrium to the right, producing more ammonia.
• Production of Calcium Oxide (Lime): In the thermal decomposition of limestone, $\text{CaCO}_3(s) \rightleftharpoons \text{CaO}(s) + \text{CO}_2(g)$, carbon dioxide gas is constantly removed from the kiln. This drives the reaction forward, ensuring the complete conversion of $\text{CaCO}_3$ to $\text{CaO}$.

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A Visual Demonstration: The Iron(III) Thiocyanate Equilibrium

This is a classic experiment that lets you see Le Chatelier's principle in action. The reaction is:

(yellow) (colourless) (deep red)

When you mix a solution of iron(III) nitrate ($\text{Fe}^{3+}$) with a solution of potassium thiocyanate ($\text{SCN}^-$), a deep red colour appears due to the formation of the complex ion $[\text{Fe(SCN)}]^{2+}$. The intensity of this red colour is directly proportional to the concentration of the complex. Once the colour intensity becomes constant, the system is at equilibrium.

You can now shift this equilibrium in either direction by changing the concentration of a reactant or product.

Shifting the Equilibrium to the Left (Reverse Direction)

You can shift the equilibrium to the left by removing one of the reactants. This is done by adding a reagent that reacts with and "ties up" either $\text{Fe}^{3+}$ or $\text{SCN}^-$ ions.

1. Removing $\text{Fe}^{3+}$: Adding oxalic acid ($\text{H}_2\text{C}_2\text{O}_4$) causes it to react with $\text{Fe}^{3+}$ ions to form a very stable complex ion, $[\text{Fe(C}_2\text{O}_4)_3]^{3-}$. This drastically reduces the concentration of free $\text{Fe}^{3+}(aq)$.

   • The Disturbance: $[\text{Fe}^{3+}]$ decreases.
   • The Response: According to Le Chatelier's principle, the system tries to replenish the lost $\text{Fe}^{3+}$. It does this by dissociating some of the $[\text{Fe(SCN)}]^{2+}$ complex. The equilibrium shifts to the left.
   • The Observable Effect: Because the concentration of the red $[\text{Fe(SCN)}]^{2+}$ decreases, the intensity of the red colour fades.

2. Removing $\text{SCN}^-$: Adding a solution of mercury(II) chloride ($\text{HgCl}_2$) causes $\text{Hg}^{2+}$ ions to react with $\text{SCN}^-$ ions to form a stable complex, $[\text{Hg(SCN)}_4]^{2-}$. This removes free $\text{SCN}^-$ from the solution.

   • The Disturbance: $[\text{SCN}^-]$ decreases. …