Equilibrium · Lesson 1 of 14
Physical Equilibrium and Dynamic Change
“See how opposing physical processes keep a system steady while particles continue to move.”
• Recognise dynamic equilibrium from equal opposing rates and constant observable properties. • Explain phase equilibria and distinguish evaporation from boiling. • Describe saturated solutions and the exchange between dissolved and undissolved material. • Use gas pressure and temperature to explain dissolved-gas behaviour.
A steady appearance can hide continuous change
A sealed bottle can contain liquid water and water vapour for a long time without an obvious change in either amount. This does not mean that molecules have stopped moving. Some leave the liquid while others return to it. Equilibrium is a condition in which these opposing transfers balance. We begin with physical processes because the transfers are easier to picture than the breaking and making of chemical bonds.
A state in which opposing processes continue at equal rates, so there is no net change in the macroscopic properties of the system under the specified conditions.
Macroscopic properties are quantities we can measure for the sample as a whole, such as pressure, concentration and the amounts of its phases. Equal rates do not require equal amounts in the two phases. For the examples here, keep temperature fixed and prevent material from escaping. A closed system can exchange energy with its surroundings while retaining its material; it need not be thermally isolated.
Melting and freezing balance at a particular temperature
Place ice and liquid water together in a well-insulated container at atmospheric pressure. At approximately 273 K, molecules leave the ordered solid and enter the liquid, while other liquid molecules join the solid. If neither phase is being used up, the rate of melting equals the rate of freezing. Insulation prevents an ongoing heat transfer from gradually melting all the ice or freezing all the water.
The normal melting point is the coexistence temperature at one atmosphere, about 1.013 bar. Changing pressure can change a melting point. The word normal specifies the pressure; it does not mean that every solid–liquid equilibrium occurs at 273 K.
Problem
A container has 20 g of ice and 80 g of liquid water at the melting temperature. Must their masses become equal?
- 1.Equilibrium compares rates of transfer, not the masses present.
- 2.If 0.2 g melts in a short interval while 0.2 g freezes, both phase masses are unchanged.
- 3.The sample may remain at 20 g of ice and 80 g of liquid. Equal masses are not required.
Vapour pressure grows until condensation catches up
Initially dry air above water contains few water molecules. Molecules with enough energy escape the liquid surface, and the water-vapour pressure rises. As more vapour accumulates, more molecules strike the liquid and remain there. Condensation becomes faster until it balances evaporation. The resulting water-vapour pressure is the equilibrium vapour pressure at that temperature, provided some liquid remains.
Heating usually raises a liquid’s vapour pressure because a larger fraction of molecules can escape. At the same temperature, a more volatile liquid has a higher vapour pressure. In open watch glasses, acetone, ethanol and water eventually evaporate because escaped vapour disperses into the room. The evaporation time depends on the liquid, its amount and temperature. These open samples do not retain enough vapour to establish the closed liquid–vapour equilibrium being studied.
Boiling occurs when a liquid’s equilibrium vapour pressure equals the external pressure, so vapour bubbles can persist within the liquid. The normal boiling point uses one atmosphere: water boils near 100 degrees Celsius at this pressure. A lower external pressure, as at high altitude, lowers the boiling temperature. Evaporation can occur below the boiling point and takes place at the surface.
Problem
A sealed liquid–vapour system at fixed temperature is suddenly expanded. Predict the immediate and eventual vapour pressure, assuming liquid remains.
- 1.Immediately after expansion, the same vapour molecules occupy a larger volume, so their pressure falls.
- 2.The lower vapour concentration reduces condensation initially. Evaporation temporarily exceeds condensation.
- 3.More liquid evaporates until the original equilibrium vapour pressure is restored at the same temperature. If all liquid evaporates first, this conclusion no longer applies.
Solids and solutions also exchange particles
Iodine in a sealed vessel produces violet vapour by sublimation, the direct transfer from solid to gas. Eventually its colour stops intensifying because deposition from vapour balances sublimation. Camphor provides another physical example. Ammonium chloride, sometimes described as subliming, involves chemical dissociation and recombination on heating; it should not be treated as an unchanged molecular vapour in the same way as iodine.
A saturated sugar solution in contact with excess sugar crystals has a fixed dissolved concentration at a specified temperature. Sugar molecules continue to leave the crystals and return to them. Adding labelled sugar crystals to an unlabelled saturated solution eventually distributes the label between solution and crystals. This tracer observation reveals exchange even when the total dissolved concentration stays unchanged. Cooling a concentrated hot syrup can produce crystals because its equilibrium solubility changes.
A solution at its equilibrium solubility at a specified temperature. When excess solute is present, dissolution and crystallisation balance.
Gas molecules also exchange between a liquid and the gas above it. For a dilute dissolved gas obeying Henry’s law at fixed temperature, its dissolved amount in a given quantity of solvent is proportional to the gas’s partial pressure. A sealed soda bottle retains much more carbon dioxide than an opened bottle because its CO2 pressure is higher. Gas escapes after opening until the dissolved amount moves towards the value appropriate to the lower CO2 partial pressure in air. For this example, warming also reduces gas solubility.
| System | Opposing transfers | Quantity fixed under stated conditions |
|---|---|---|
| Ice and water | Melting and freezing | Coexistence temperature at fixed pressure |
| Liquid and vapour | Evaporation and condensation | Vapour pressure at fixed temperature |
| Solid and vapour | Sublimation and deposition | Vapour pressure at fixed temperature |
| Solid and saturated solution | Dissolution and crystallisation | Dissolved concentration at fixed temperature |
| Gas and its solution | Dissolution and escape | Dissolved concentration / gas partial pressure at fixed temperature |
Problem
The dissolved concentration of a gas is 0.020 M at a partial pressure of 2.0 bar. Estimate it at 0.50 bar at the same temperature, assuming Henry’s law applies.
- 1.At fixed temperature the ratio of dissolved concentration to partial pressure is constant.
- 2.The pressure ratio is .
- 3.The new dissolved concentration is . The gas becomes less soluble as its partial pressure falls.
A constant pressure or concentration is evidence to interpret, not proof that particles have stopped. Check whether both opposing processes remain possible and whether the required phases are still present.
Quiz
What is equal at dynamic equilibrium?
At fixed temperature, more liquid water is added to a sealed liquid–vapour equilibrium. Some liquid was already present. What happens to its final vapour pressure?
Why does soda release dissolved carbon dioxide when opened?
Why does radioactive sugar appear in both crystals and saturated solution?
A liquid boils at a lower temperature on a mountain because
Which condition is needed for the final vapour pressure after expansion to return to its original value?
Practice Problems
- Explain why 10 g of ice can coexist with 90 g of water without becoming equal in mass.
- Describe the observations in the sealed vapour-pressure apparatus from the beginning to equilibrium.
- Compare evaporation from open watch glasses of water and acetone at the same temperature.
- Explain why sugar crystallises when concentrated hot syrup cools.
- A gas has dissolved concentration 0.030 M at 3.0 bar. Find it at 1.2 bar at the same temperature under Henry-law conditions.
- Explain what labelled sugar reveals that a concentration measurement alone cannot.
Key Takeaways
Physical equilibrium balances opposing transfers; it does not stop particle motion. Temperature, pressure and the phases present determine the equilibrium condition. Vapour pressure, solubility and gas dissolution are observable expressions of molecular exchange.
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How Chemical Equilibrium Is Established