# Static and Dynamic Equilibrium

> Reversible reactions, open and closed systems, dynamic and static equilibrium, and how enthalpy, entropy and Gibbs free energy decide if a reaction happens by itself.

From HSC Chemistry on WhizzNotes, free study notes for NSW students: https://whizznotes.com/chemistry/static-and-dynamic-equilibrium/. Updated 3 November 2026.
Year 11 and Year 12 boxes go deeper than the Year 10 text around them.

### Story time: The busy escalator
Mia stands at the top of a shopping centre with two escalators: one going up, one going down. People ride both all afternoon. Every minute, about 30 people go up and 30 people come down.

From the balcony, the number of shoppers on the top floor never seems to change. But nobody is standing still. People are moving both ways at the same speed, so the totals stay the same.

A chemical reaction in a sealed flask can do exactly this. Molecules keep reacting forwards and backwards, but because both happen at the same rate, nothing seems to change. Chemists call this **dynamic equilibrium**.

**In short:** In a reversible reaction in a closed system, equilibrium is reached when the forward and reverse reactions go at the same rate. This is dynamic equilibrium: both reactions keep going, but the amounts of each substance stay the same. In static equilibrium, all reactions have stopped.

## Reversible reactions

Most reactions you met in Year 10 go one way. Burn a match and you cannot turn the ash and smoke back into a match.

A **reversible reaction** can go both ways. We write it with a double arrow:

$$\mathrm{N_2O_4(g)} \rightleftharpoons \mathrm{2NO_2(g)}$$

- The **forward reaction** goes left to right.
- The **reverse reaction** goes right to left.

Colourless $\mathrm{N_2O_4}$ breaks into brown $\mathrm{NO_2}$, and brown $\mathrm{NO_2}$ joins back up into $\mathrm{N_2O_4}$, at the same time.

## Open and closed systems

- An **open system** swaps both **matter and energy** with its surroundings. An open beaker of fizzy drink loses carbon dioxide gas into the air.
- A **closed system** swaps **energy but not matter**. A sealed bottle keeps all its gas inside but can still warm up or cool down.

**Equilibrium can only be reached in a closed system.** If a product escapes (like a gas bubbling out of an open beaker), the reverse reaction cannot happen, so the reaction just runs to the end.

## Dynamic equilibrium

Put $\mathrm{N_2O_4}$ in a sealed flask and watch what happens:

1. At the start, only the forward reaction can happen, and it is fast because there is lots of $\mathrm{N_2O_4}$.
2. As $\mathrm{N_2O_4}$ is used up, the forward reaction slows down.
3. As $\mathrm{NO_2}$ builds up, the reverse reaction speeds up.
4. Soon the two rates become **equal**. From then on, the amounts of each gas stay the same.

This state is **dynamic equilibrium**:

- The forward and reverse reactions are **both still happening**.
- Their **rates are equal**.
- The **concentrations stay constant** (but are usually not equal to each other).
- **Macroscopic properties** (colour, pressure, temperature) stay constant.

| | Static equilibrium | Dynamic equilibrium |
|---|---|---|
| Are reactions happening? | No, they have stopped | Yes, both directions |
| Rates | Both zero | Forward = reverse, not zero |
| Example | A fully burnt-out fire | $\mathrm{N_2O_4}$ and $\mathrm{NO_2}$ in a sealed flask |

In **static equilibrium**, nothing is reacting at all. A finished combustion, or diamond slowly "turning into" graphite at room temperature (far too slow to notice), is static.

### Worked example: Reading the graphs
A **concentration-time graph** shows the reactant line falling and the product line rising, until both lines go flat. Flat lines mean equilibrium.

A **rate-time graph** shows the forward rate starting high and falling, the reverse rate starting at zero and rising. Where the two lines meet and run together, equilibrium has been reached.

## Non-equilibrium systems

Some reactions are reversible on paper but never reach equilibrium in real life:

- **Combustion** gives out lots of energy and the gases escape. It is effectively one way.
- **Photosynthesis** and **respiration** are the reverse of each other, but they happen in different places, with different enzymes, so they are not an equilibrium.

Heating blue hydrated copper(II) sulfate turns it white as the water leaves. Adding water turns it blue again. This shows the reaction is **reversible**, but because the water escapes as steam in an open tube, it is **not** an equilibrium. Cobalt(II) chloride crystals (pink when hydrated, blue when dry) show the same thing.

## Enthalpy and entropy

Two things decide whether a reaction happens by itself.

**Enthalpy change** $\Delta H$ is the heat energy taken in or given out.

- **Exothermic**: $\Delta H$ is negative. Heat is given out. Nature favours this.
- **Endothermic**: $\Delta H$ is positive. Heat is taken in.

**Entropy** $S$ measures how spread out (disordered) energy and particles are. Nature favours **more disorder**, so a positive entropy change $\Delta S$ helps a reaction happen.

Entropy goes **up** ($\Delta S > 0$) when:

- a solid melts, or a liquid boils (solid < liquid < gas),
- a solid dissolves,
- a reaction makes **more moles of gas** than it uses.

## Gibbs free energy

**Gibbs free energy** combines both ideas:

$$\Delta G = \Delta H - T\Delta S$$

where $T$ is the temperature in **kelvin** ($T = °\text{C} + 273$).

- $\Delta G < 0$: the reaction is **spontaneous** (it can happen by itself).
- $\Delta G > 0$: it is not spontaneous. It needs a steady supply of energy.
- $\Delta G = 0$: the system is at equilibrium.

| $\Delta H$ | $\Delta S$ | Spontaneous? |
|---|---|---|
| − | + | Always |
| + | − | Never |
| − | − | Only at low temperature |
| + | + | Only at high temperature |

**Watch the units.** $\Delta H$ is usually in kJ, but $\Delta S$ is in J/K. Divide $\Delta S$ by 1000 before you use the formula.

### Worked example: Is it spontaneous at 25 °C?
For a reaction, $\Delta H = -92$ kJ and $\Delta S = -199$ J/K.

$T = 298$ K and $\Delta S = -0.199$ kJ/K.

$\Delta G = -92 - 298 \times (-0.199) = -92 + 59.3 = -32.7$ kJ

$\Delta G$ is negative, so the reaction is spontaneous at 25 °C.

Photosynthesis has $\Delta H > 0$ and $\Delta S < 0$ (six gas molecules become one solid sugar and six gas molecules, and energy is stored). So $\Delta G$ is always positive. That is why plants need a constant supply of **light energy**.

**Mind map**

```text
Equilibrium
  Reversible reactions
    Double arrow
    Forward and reverse
  Systems
    Open: matter and energy escape
    Closed: energy only
    Equilibrium needs closed
  Dynamic
    Rates equal, not zero
    Concentrations constant
    Properties constant
  Static
    Reactions stopped
  Spontaneity
    ΔH negative helps
    ΔS positive helps
    ΔG = ΔH − TΔS
    ΔG < 0 spontaneous
```

### Year 11: Energy profiles and catalysts
Particles only react when they **collide** with enough energy, the **activation energy** $E_a$, and the right orientation. This is **collision theory**.

On an energy profile diagram:

- $E_a$ (forward) is the climb from the reactants to the top of the hill.
- $\Delta H$ is the drop (or rise) from reactants to products.
- $E_a$ (reverse) $= E_a$ (forward) $- \Delta H$.

For example, if $E_a$ (forward) $= 120$ kJ and $\Delta H = -40$ kJ, then $E_a$ (reverse) $= 120 + 40 = 160$ kJ.

A **catalyst** gives a new path with a lower hill. It lowers $E_a$ for the forward and reverse reactions **by the same amount**, so both speed up equally. It does **not** change $\Delta H$, and it does **not** change the amounts at equilibrium. Equilibrium is just reached sooner.

### Year 12: The turning temperature
When $\Delta H$ and $\Delta S$ have the same sign, the reaction changes from not spontaneous to spontaneous at the temperature where $\Delta G = 0$:

$$T = \frac{\Delta H}{\Delta S}$$

**Example.** Heating limestone: $\mathrm{CaCO_3(s)} \rightarrow \mathrm{CaO(s)} + \mathrm{CO_2(g)}$, with $\Delta H = +178$ kJ and $\Delta S = +161$ J/K.

$T = \dfrac{178\,000}{161} = 1106$ K, or about 833 °C. Above this, $\Delta G < 0$ and the limestone breaks down by itself.

**Entropy change from tables.** $\Delta S = \sum S(\text{products}) - \sum S(\text{reactants})$, multiplying each value by its number in the equation. For $\mathrm{2SO_2} + \mathrm{O_2} \rightarrow \mathrm{2SO_3}$ with $S$ values 248, 205 and 257 J/K/mol:

$\Delta S = 2(257) - [2(248) + 205] = 514 - 701 = -187$ J/K. Three gas molecules become two, so entropy falls, as expected.

**Below or above?** If $\Delta H$ and $\Delta S$ are both negative, the reaction is spontaneous **below** $T = \dfrac{\Delta H}{\Delta S}$. For $\Delta H = -50.0$ kJ and $\Delta S = -150$ J/K, that is below 333 K.

**Reading a graph of $\Delta G$ against $T$.** $\Delta G = \Delta H - T\Delta S$ is a straight line. Its intercept at $T = 0$ is $\Delta H$ and its gradient is $-\Delta S$. A line falling from positive values means $\Delta S > 0$, and where it crosses zero the reaction becomes spontaneous.

**Two values of $\Delta G$.** If $\Delta G = -20$ kJ at 300 K and $-40$ kJ at 500 K, then $\Delta S = \dfrac{-20 - (-40)}{500 - 300} = 0.10$ kJ/K, and $\Delta H = -20 + 300 \times 0.10 = +10$ kJ.

**Concentrations from a graph.** Use the mole ratio. In $\mathrm{A} \rightleftharpoons \mathrm{2B}$, if $[\mathrm{A}]$ falls by 0.10 mol/L, $[\mathrm{B}]$ rises by 0.20 mol/L.

**Heat from a fuel.** The heat needed is $q = mc\Delta T$. Heating 2.0 kg of water by 60 °C needs $2.0 \times 4.18 \times 60 = 502$ kJ. Carbon gives 394 kJ per mole, so $\dfrac{502}{394} = 1.27$ mol, or 15.3 g of carbon.

## Quick summary

- A reversible reaction goes both ways. Equilibrium needs a **closed system**.
- **Dynamic equilibrium**: forward rate = reverse rate (not zero), concentrations and properties stay constant.
- **Static equilibrium**: all reactions have stopped.
- Combustion and photosynthesis/respiration are not equilibrium systems.
- $\Delta H < 0$ (exothermic) and $\Delta S > 0$ (more disorder) both favour a reaction.
- $\Delta G = \Delta H - T\Delta S$, with $T$ in kelvin and $\Delta S$ in kJ/K. $\Delta G < 0$ means spontaneous.

## Quiz

Try these questions from the chapter quiz.

1. In dynamic equilibrium:
   - A. Both reactions have stopped
   - B. The forward and reverse rates are equal
   - C. The concentrations of reactants and products are equal
   - D. Only the forward reaction happens
   - Answer: B. The forward and reverse rates are equal. Both reactions keep going at the same rate, so the amounts stay constant.

2. Equilibrium can only be reached in:
   - A. An open system
   - B. A closed system
   - C. A system with no reactants
   - D. A system that is being heated
   - Answer: B. A closed system. If matter can escape, the reverse reaction cannot keep up.

3. A closed system can exchange:
   - A. Matter only
   - B. Energy only
   - C. Both matter and energy
   - D. Neither matter nor energy
   - Answer: B. Energy only. A sealed flask can warm up or cool down but keeps its chemicals inside.

The whole chapter quiz is on the page, for signed-in students: https://whizznotes.com/chemistry/static-and-dynamic-equilibrium/#quiz