Solutions Sugar or alcohol in water A homogeneous mixture in which all of the particles have the sizes of atoms. Driving forces for solution formation Spontaneous tendency for increasing disorder (entropy!) 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.

Solution formation

Alcohol (ethanol) in water

Engine coolants We can now explain why car radiator coolants dissolve in water. The coolants typically contain either ethylene glycol or propylene glycol, which, like ethanol and water, contain hydrogen-bonding O-H bonds.

Solublility & Miscibility (like dissolves like)

How do stain removers work? Like dissolves like dry cleaning, surfactants and oxidizing agents

Energy Changes and Solution Formation Heat of solution: The enthalpy change between system and surroundings when 1 mole of a solute dissolves in solvent at constant pressure - In general, solutions form when Hsoln is negative. When Hsoln is too positive, a solution will not form. Heat of solution is zero for an ideal solution Examples: MgSO4 added to water has Hsoln = -91.2 kJ/mol. NH4NO3 added to water has Hsoln = + 26.4 kJ/mol. MgSO4 is used in hot packs and NH4NO3 is used in cold packs.

Ideal dilute solutions-effect of pressure on solubility PB = xBKB

Pressure and Solubility “The Bends” - deep sea divers N2 in compressed air has low solubility in blood. At great depths, partial pressure of N2 increases and solubility increases. After returning to surface, solubility decreases substantially and N2 come out of blood causing small bubbles in the capillaries which may rupture the blood vessels - fatal - risk reduced if He is used since it has lower Henry’s constant. Soft drinks and champagne produced by dissolving CO2 in a liquid under pressure opening bottle reduces the partial pressure of CO2 above solution, solubility reduced, CO2 leaves solution (effervescence)

Molecular interpretation of Raoult’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 (see later): when a solute is added to a pure liquid, the entropy (disorder) 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. The effect shifts the s-l and l-g vp curve of a phase diagram down.

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

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