Chemical Equilibrium || Le Chatelier's Principle || Lecture # 2 || Dr. Rizwana

TL;DR

Kc changes with temperature because Kc is linked to ΔH; exothermic reactions (ΔH < 0) show decreasing Kc when temperature rises, while endothermic reactions (ΔH > 0) show increasing Kc when temperature rises.

Briefing Cornell Notes

Briefing

Chemical equilibrium constant (Kc) stays fixed against concentration, pressure, and catalyst changes only in specific ways—but temperature changes reliably shifts Kc. The lecture frames this through Le Chatelier’s principle: when a system at equilibrium is disturbed, the reaction shifts to counter the disturbance. Concentration and pressure adjustments change how much the system “wants” to react in the forward versus reverse direction, yet the equilibrium constant’s value remains tied to temperature, not to those other physical tweaks.

Temperature is the key exception. For an equilibrium reaction, Kc depends on temperature through the enthalpy change (ΔH). When ΔH is negative (an exothermic, “heat-evolving” reaction), increasing temperature drives Kc downward. The reason is practical: added heat effectively accelerates the backward (reverse) direction, pushing the system to consume products and reform reactants. When ΔH is positive (an endothermic, “heat-absorbing” reaction), increasing temperature raises Kc because the added thermal energy favors the forward direction, increasing the equilibrium amount of products.

Pressure changes are treated next using the ammonia synthesis example: nitrogen and hydrogen form ammonia, with stoichiometry that matters for gas-phase equilibria. The lecture emphasizes that applying pressure changes the effective “packing” of gas molecules in a confined volume. Le Chatelier’s principle then predicts a shift toward the side with fewer moles of gas, because that direction reduces the stress caused by compression. In the ammonia example, pressure favors the formation of ammonia when the forward reaction reduces the number of gas molecules relative to the reactant side. The equilibrium composition adjusts accordingly, but the equilibrium constant’s value is not presented as changing—rather, the system’s direction and rates adjust to minimize the applied pressure effect.

Concentration changes follow the same logic: if extra reactant or product is added, the system shifts to consume the added species and restore equilibrium. Using the nitrogen–hydrogen–ammonia setup again, adding more reactant (like hydrogen) pushes the reaction forward until the added concentration is consumed; removing product (like ammonia) similarly drives the reaction forward to replenish it. The lecture stresses that equilibrium is re-established only after the system consumes the “excess” concentration, meaning the forward and reverse rates rebalance again.

Catalysts are the final physical parameter. Catalysts do not change Kc. They lower activation energy, speeding up both forward and reverse reactions by providing an easier pathway, so equilibrium is reached faster but at the same equilibrium constant value. In short: temperature shifts Kc through ΔH; concentration and pressure shift equilibrium position to counter the disturbance; catalysts accelerate the approach to equilibrium without changing the equilibrium constant.

Cornell Notes

The equilibrium constant Kc depends on temperature and the reaction’s enthalpy change (ΔH), not on concentration, pressure, or catalysts. Increasing temperature decreases Kc for exothermic reactions (ΔH < 0) because it favors the reverse direction, while it increases Kc for endothermic reactions (ΔH > 0) by favoring the forward direction. Changing pressure shifts the equilibrium position in the gas-phase direction that reduces the number of moles of gas, illustrated with ammonia formation from nitrogen and hydrogen. Changing concentration shifts the reaction to consume added reactants or replace removed products until equilibrium is restored. Catalysts lower activation energy and speed up both directions equally, so Kc stays the same.

Why does changing temperature alter Kc, while changing concentration or pressure mainly shifts equilibrium position?

Kc is tied to temperature through the reaction enthalpy change (ΔH). When temperature increases, the system responds by shifting toward the direction that counteracts the added heat. For ΔH < 0 (exothermic), added heat favors the reverse reaction, so Kc decreases. For ΔH > 0 (endothermic), added heat favors the forward reaction, so Kc increases. By contrast, concentration and pressure disturbances change which side is favored at equilibrium, but the lecture treats Kc as not being directly changed—equilibrium is re-established with a new composition while Kc remains governed by temperature.

How does Le Chatelier’s principle predict the effect of pressure on ammonia synthesis?

Pressure changes the effective “packing” of gas molecules in a confined volume. Le Chatelier’s principle says the system shifts to minimize the applied stress, which for gases means shifting toward the side with fewer moles of gas. In the ammonia example, nitrogen and hydrogen react to form ammonia with stoichiometry that makes the gas-mole count differ between reactants and products. Applying pressure favors the direction that reduces the number of gas molecules, so ammonia formation is promoted under compression.

What happens when extra reactant is added to a system already at equilibrium?

Adding extra reactant increases its concentration at the equilibrium stage. Le Chatelier’s principle predicts the reaction will proceed in the direction that consumes the added species. In the nitrogen–hydrogen–ammonia example, adding more hydrogen pushes the reaction forward to produce more ammonia until the added hydrogen is consumed and the system returns to equilibrium. Equilibrium persists only once the concentrations match the equilibrium condition again.

What happens if product is removed from an equilibrium mixture?

Removing product lowers its concentration below the equilibrium value. The system responds by shifting toward the side that restores the removed product. In the ammonia example, removing ammonia drives the reaction forward so nitrogen and hydrogen combine again to replenish ammonia, until equilibrium is re-established.

Why doesn’t a catalyst change Kc even though it speeds up reactions?

A catalyst lowers activation energy for both forward and reverse reactions. That speeds up how quickly the system reaches equilibrium, but it does not change the equilibrium constant Kc. Since Kc depends on temperature (and ΔH), not on the activation-energy barrier, the equilibrium composition at a given temperature remains the same—only the time to reach it decreases.

Review Questions

If a reaction has ΔH < 0, what direction does increasing temperature push the equilibrium, and how should Kc change?
In a gas-phase equilibrium, how does the number of moles of gas on each side determine the shift caused by increasing pressure?
Why does adding a catalyst speed up reaching equilibrium without changing Kc?

Key Points

1
Kc changes with temperature because Kc is linked to ΔH; exothermic reactions (ΔH < 0) show decreasing Kc when temperature rises, while endothermic reactions (ΔH > 0) show increasing Kc when temperature rises.
2
Changing concentration shifts the equilibrium position to consume added reactants or replace removed products until equilibrium is restored.
3
Applying pressure shifts gas-phase equilibria toward the side with fewer moles of gas to reduce the stress caused by compression.
4
For pressure effects, the equilibrium composition changes, but Kc is treated as remaining governed by temperature.
5
Catalysts do not change Kc because they lower activation energy for both forward and reverse reactions equally.
6
Catalysts reduce the time required to reach equilibrium without altering the equilibrium constant value at a fixed temperature.

Highlights

Temperature is the decisive lever for Kc: increasing temperature decreases Kc for exothermic reactions and increases Kc for endothermic ones.

Pressure shifts equilibrium toward the side with fewer gas moles, illustrated through ammonia formation from nitrogen and hydrogen.

Concentration changes drive the system to consume added species or replace removed products until equilibrium is re-established.

Catalysts speed up both directions by lowering activation energy but leave Kc unchanged.

Topics

Chemical Equilibrium
Le Chatelier’s Principle
Equilibrium Constant
Temperature Dependence
Catalysts