Vapour Pressure and Heat
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.
Phase changes can be expressed as
enthalpy changes at constant temperatures
(Claussius-Clapeyron equation).
vapH


ln P o  
 P 
R
 1 1
 o 
T
T
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 fluids
Heating 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 diagrams
Water
Carbon dioxide
Triple pt: 6.11mbar, 273.16K
Triple pt: 5.11bar, 218.8K
Critical pt: 215bar, 647.3K
Critical pt: 72bar, 304.2K
Solutions
A 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 water
Glucose 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 Solution
The enthalpy change between system and surroundings when
1 mole of a solute dissolves in solvent at constant pressure
Vapourised particles and solvent
solvation energy
-lattice energy
solnH
solid
solution
solvent
Heat of solution is zero
for an ideal solution
Ideal dilute solutions
PB = 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 mixtures
Deviations 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.
T f  K f m
Tb  K b m
Osmotic pressure
Osmotic membrane: semi-permeable membrane
that allows passage of only solvent molecules
Dialysis membrane: membrane that allows
passage of solvent and small solutes.
Van’t Hoff equation
  MRT
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+ + ClWe 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
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