How far? The extent of chemical change

A reaction that can go in **both directions** — reactants can form products and products can re-form reactants. Shown with the **⇌** symbol.

The state, in a **closed system**, where the **forward and reverse reactions occur at equal rates**, so the **concentrations of reactants and products stay constant**.

Because **both** the forward and reverse reactions are **still happening** — the reaction has **not** stopped; the opposite changes simply cancel out.

**No** — they are **constant** (unchanging), but generally **not equal** to one another.

The **rate of the forward reaction = the rate of the reverse reaction**.

So that **nothing is added or escapes** (no reactant/product/heat leaves); an open system could never settle to constant concentrations.

**Colour (absorbance)**, **pressure** (for gases) and **pH** — all remain constant because the concentrations are constant.

**How far** a reaction has gone — the **relative amounts** of reactants and products. To the **right** = mostly products; to the **left** = mostly reactants.

Measure a **macroscopic property** over time (e.g. colour intensity); when it **levels off to a constant value**, equilibrium has been reached.

That the reaction has **stopped** — in fact both reactions continue (it is **dynamic**); the concentrations are merely constant.

The **forward rate falls** and the **reverse rate rises** until they **meet (become equal)** — that point is equilibrium.

Both the reactant and product curves **level off** (become flat) and stay constant — at **different** values.

If a system at equilibrium is disturbed, the position shifts in the direction that **opposes** (partly cancels) the change.

…towards the **products** (right) — the system uses up the added reactant.

…towards the **products** (right) — the system replaces the lost product.

The position shifts to the side with **fewer moles of gas**, to reduce the pressure.

Changing the pressure causes **no shift** in the position.

The position shifts in the **endothermic** direction (it absorbs the added heat).

A change in **temperature** — concentration, pressure and a catalyst leave K_{c} unchanged.

**No shift** in position and **no change** in K_{c}; it just reaches equilibrium **sooner** (speeds up both directions equally).

K_{c} **decreases** (the position shifts towards the reactants).

**Count the moles of gas** on each side; the position shifts towards the side with **fewer** gas moles when pressure rises.

Write **heat** as a species (exo: products + heat; endo: reactants + heat), then treat adding heat like adding that species.

**Endothermic** — adding heat favours the heat-absorbing direction.

The **fixed ratio** of product to reactant concentrations at equilibrium, at a given temperature — it shows **how far** a reaction goes.

**Products over reactants**, each concentration **raised to the power of its balancing coefficient**; use [ ] for equilibrium concentration in mol dm⁻³.

$K_{c} = \dfrac{[\text{NH}_{3}]^{2}}{[\text{N}_{2}][\text{H}_{2}]^{3}}$ — the 2 and 3 become powers.

**Products are favoured** — the equilibrium lies to the **right** (mostly products).

**Reactants are favoured** — the equilibrium lies to the **left** (mostly reactants).

The **reciprocal**: K_{reverse} = **1 / K_{forward}**.

A change in **temperature** — concentration, pressure and a catalyst all leave K_{c} unchanged.

K_{c} **increases** (the position shifts towards products).

K_{c} **decreases** (the position shifts towards reactants).

Convert each **amount (mol)** to a **concentration (mol dm⁻³)** using **c = n/V**, then substitute.

Pure **solids** and pure **liquids** — only gases and dissolved (aqueous) species appear.

**No** — K_{c} describes the **extent** (how far), not the **rate** (how fast).