Vapour Pressure and Heat

Phase changes can be expressed as enthalpy changes at constant

temperatures (Claussius-Clapeyron equation).

What happens to a solid substance when it is heated?

The compound can simply get hotter or a phase change can occur.

The transition from the solid phase to the liquid phase is an example of a phase change, which is often called melting. Boiling or vapourisation is an example of a phase change from the liquid to the gas phase.

TTR

H

PP

o

vapo

11ln

Phase diagrams

Show regions of P and T at which various phases are thermodynamically stable

Triple point: three phases in equilibrium

Vp curve for solid

Vp curve for liquid

Melting point line

Critical point

Supercritical fluidsHeating liquid in a closed vessel does not produce boiling. The vp (density) of the vapour rises with increasing T, as the liquid density decreases, until both are equal and a single phase exists (neither liq nor vap.).

Typical phase diagramsWater Carbon dioxide

Triple pt: 6.11mbar, 273.16K

Critical pt: 215bar, 647.3K

Triple pt: 5.11bar, 218.8K

Critical pt: 72bar, 304.2K

SolutionsA homogeneous mixture in which all of the particles have the sizes of atoms.

Driving forces for solution formation

(i) Spontaneous tendency for increasing disorder (entropy!)

(ii) Intermolecular forces

Sugar or alcohol in waterGlucose has -O-H groups along the carbon skeleton. These -O-H are polar centers.

Glucose dissolves in water because polar water molecules attach to the glucose molecules by dipole-dipole (H-bond) forces. When the attractive forces of the water molecules for the glucose exceeds the attractive forces between the glucose and its neighbouring glucose molecules the water can rip the sugar molecule out of the crystal. The glucose is "solvated" when it surrounded solvent molecules. The solvent has "dissolved" the molecule.

Water and ethyl alcohol are completely "miscible". Both water and ethanol are polar molecules with hydrogen bonding. The similarity of the two molecules results in solutions where the water and alcohol molecules are interchangeable.

Heats of SolutionThe enthalpy change between system and surroundings when

1 mole of a solute dissolves in solvent at constant pressure

solid solvent

Vapourised particles and solvent

solution

-lattice energy solvation energy

solnH

Heat of solution is zero for an ideal solution

Ideal dilute solutionsPB = xBKB Henry’s Law

Colligative Properties of Solutions

Physical properties that depend only upon the populations of particles in a mixture

Effect of solutes on the vapour pressure of solutions

Psoln = xsolventP*solvent Raoult’s Law

Molecular interpretation of Raoult’s Law

Volatile solutes and Raoult’s Law

Each component contributes its own partial pressure to the

solution vapour pressure (Dalton’s Law)

Real mixturesDeviations because of

intermolecular attractions

Boiling point elevation Freezing point depression

Entropy effect: when a solute is added to a pure liquid, the entropy is increased relative to the vapour phase. Therefore there is a weaker tendency to form a vapour (boiling point elevation). A similar molecular interpretation explains freezing point depression.

mKT

mKT

bb

ff

Osmotic pressureOsmotic membrane: semi-permeable membrane that allows passage of only solvent molecules

Dialysis membrane: membrane that allows passage of solvent and small solutes.

MRTVan’t Hoff equation

Colligative properties of solutions of electrolytes

1.00 m NaCl: F.P= -3.37C (not –1.86C as expected)!

Colligative properties depend on the concentration of particles

Remember: NaCl Na+ + Cl-

We have 2.00m of particles and should get F.P: -(2x1.86C) = -3.72C

Effect of interionic attractions account for discrepancy between actual and calculated F.P. for ionic species.

Van’t Hoff Factor compares degrees of dissociation of electrolytes

calculated

measured

T

Ti