Chemical Equilibrium Lab

Most reactions are reversible. They reach a state where forward and reverse rates are equal — equilibrium. It looks static from the outside, but both reactions are still running flat out; they just cancel each other out.

The equilibrium constant

For a reaction aA + bB ⇌ cC + dD:

K = [C]^c[D]^d / [A]^a[B]^b

K is fixed for a given temperature. If K is large, products dominate; if small, reactants do. Crucially, only temperature changes K — changing concentrations or pressure shifts the position but not the constant itself.

Le Chatelier’s principle

“If a system at equilibrium is disturbed, it shifts to counteract the disturbance.” Concretely:

• Add a reactant → the system consumes it, making more product.
• Remove a product → the system makes more to replace it.
• Raise the pressure (gases) → the system favours the side with fewer gas molecules.
• Raise the temperature → the system favours the endothermic direction, absorbing the added heat.

Temperature is the subtle one: treat heat as a reactant for an endothermic reaction and as a product for an exothermic one, and the rule becomes intuitive.

Try this

• Add reactant and watch product climb, then settle at a new position with the same K.
• For an exothermic reaction, raise the temperature — yield falls. This is the central tension in the Haber process: low temperature favours ammonia but makes the reaction slow, so industry compromises with a catalyst and moderate heat.
• Compress a gaseous equilibrium and count molecules on each side to predict the shift before you see it.

Why it matters

Equilibrium control is how ammonia is made for fertiliser (feeding roughly half the world), how sulphur dioxide is converted to sulphuric acid, how your blood buffers its pH, and how oxygen loads onto haemoglobin in the lungs and unloads in tissue.