General Chemistry - Properties of Solutions

Learning Outcomes

• Understand how enthalpy and entropy influence solution formation.

• Explain the relationship between intermolecular forces and solubility.

• Describe the effect of temperature on the solubility of solids and gases.

• Describe the relationship between the partial pressure of a gas and its solubility (Henry’s Law).

• Calculate the vapor pressure of a solvent over a solution.

• Calculate solute concentration using molarity, molality, mole fraction, weight % composition, and parts per million and be able to interconvert between units.

• Describe colligative properties of a solution and explain the difference between the effects of nonelectrolytes and electrolytes on the colligative properties.

• Calculate the boiling point elevation and freezing point depression.

• Calculate the osmotic pressure of a solution.

Solutions

• Solution – homogeneous mixture of two or more pure substances.

• The solute is dispersed uniformly in the solvent.

• Can be mixtures of gases, liquids, solids.

• The potential to mix substances to form a solution depends on their natural tendency of mixing and ability to overcome intermolecular forces.

Intermolecular Attraction in Solutions

• Must be overcome to allow for a homogeneous solution.

• Solute-solute interactions.

• Solvent-solvent interactions.

• New intermolecular forces available.

• Solvent-solute interactions.

Solution Formation

• Solvent-solute interactions between water molecules and NaCl allow solid to dissolve.

• Ions hydrated in solution.

Ionic Solubility

• Solubility of salts generally increases with temperature.

• Salts generally dissolve in polar substances.

• Dielectric constant – the potential of a substance to decrease the electrostatic interactions between two substances.

  • Crown ethers

  • D (dipole moment): 2.88 (for one substance), 1.85 (for another)

  • ε (dielectric constant): 20.7, 80.1

  • LiCl solubility at 20 ℃: 4.1 g, 83 g

Organic Compounds in Water

• Polar organic molecules dissolve more readily than nonpolar molecules.

• Those containing OH, COOH, NH2 offer hydrogen bonding potential.

Biological Molecules

• Fat-soluble molecules are stored; only minimal amounts are generally needed.

• Water-soluble molecules need large amounts as they are washed out readily.

Intermolecular Attraction in Solutions

• Solvent and solute must be separated into their component parts.

• \Delta H{soln} = \Delta H{solute} + \Delta H{solvent} + \Delta H{mix}

• \Delta H{mix} > \Delta H{solute} + \Delta H_{solvent} leads to a spontaneous mixing.

• \Delta H{mix} < \Delta H{solute} + \Delta H_{solvent} leads to a slow mixing that might not even happen.

Mixing

• There is a natural tendency toward mixing, which is spontaneous.

• Gases act as if alone when mixing in a closed system.

• For mixing in condensed phases to occur, intermolecular forces must be overcome.

• Mixing causes randomness which leads to an increase in entropy.

• Entropy – measure of the tendency to spread reducing capacity to do work.

• Reflects the randomness or disorder of the system.

Enthalpy of Mixing

• \Delta H{soln} = \Delta H{solute} + \Delta H{solvent} + \Delta H{mix}

Solubility

• Solubility is affected by:

  • Solute-solvent interactions

  • Pressure, especially for gaseous solutes

  • Temperature

• Over-generalization of “like dissolves like”.

• The stronger the solute-solvent interactions, the greater the solubility (miscibility).

• These gases display what type of intermolecular interactions?

Opposing Processes

• Dynamic equilibrium occurs in saturated solutions.

• A solution is saturated once no more solute can dissolve.

• Saturation is dependent on solubility and correlated with temperature.

• Supersaturated solutions have more solute than normally possible at a given temperature.

• Unstable; crystallization can be rapidly induced with a seed/scratch.

Pressure and Solubility

• Intermolecular forces for mixing solids and liquids are not appreciably affected by external pressure.

• Gas solubility is highly affected by pressure.

• Henry’s Law: S{gas} = kH P_g

  Gas	k_H
  He	3.9
  Ne	4.7
  Ar	15
  N2	7.1
  O2	14
  CO2	392

Temperature and Solubility

• Solubility of gases in liquids decreases as temperature increases.

• For most solids, solubility increases with temperature.

Aqueous Solution; Chemical Rxns

• Ni(s) + 2 HCl(aq) → NiCl2(aq) + H2(g)

• The solute is the salt.

• Evaporation of the solvent leaves some water in the crystal lattice.

• For sugar dissolution and evaporation, only rock candy is recovered.

Units of Concentration

• Mass or mol percentage.

• Parts per million (ppm) or billion (ppb).

• Mole fraction.

• Molarity.

• Molality.

Units of Concentration

• Mol fraction (χ):

  • Ratio of the mol of the solute to the total mols in the solution.

• Mass or mol percentage.

• Parts per million (ppm) or billion (ppb).

• Molarity.

• Molality.

Units of Concentration

• Mole fraction.

• Mass or mol percentage – “out of 100”:

  • Ratio of the mass/mol of the solute to the total solution.

  • Multiply by 100 to make it a %.

• Parts per million (ppm) or billion (ppb).

• Molarity.

• Molality.

Units of Concentration

• Mole fraction.

• Mass or mol percentage.

• Parts per million (ppm) or billion (ppb):

  • Ratio of the mass of a component to the solution.

  • Multiply by 10^6 (million) or 10^9 (billion).

  • Useful for very dilute systems.

  • As – 10 ppb, F – 400 ppb; Pb – 15 ppb; Hg – 2 ppb; Cd – 2 ppm.

• Molarity.

• Molality.

Units of Concentration

• Mole fraction.

• Mass or mol percentage.

• Parts per million (ppm) or billion (ppb).

• Molarity (M):

  • Moles of solute in liters of solution.

• Molality (m):

  • Moles of solute in kg of solvent.

  • For dilute solutions, molarity and molality are similar.

  • Molality does not vary with temperature.

Colligative Properties

• Solutes can change the physical properties of solvents.

  • Vapor-pressure lowering

  • Boiling-point elevation

  • Freezing-point depression

  • Osmotic pressure

• Colligative properties depend on solute quantity, not identity.

Vapor Pressure

• Adding nonvolatile solutes makes it harder for solvent to leave the surface as vapor.

• Increasing concentration of nonvolatile solutes lowers the vapor pressure of the solution compared to the pure solvent.

• Boiling point of the solution scales with concentration.

Raoult’s Law

• Vapor pressure of a volatile solvent in a solution is a mol fraction of the pure solvent.

• This law is based on the assumption of an ideal solution.

• Compounds that dissociate result in greater depression of solutions volatile solvent.

• If we have 15 mol% NaClaq solution what will the new vapor pressure of water be at 20 ˚C? (P˚_{solvent} = 17.5 torr)

  • P_{solution} = (0.85)*(17.5 \text{ torr}) = 14.9 \text{ torr} (WRONG)

  • NaCl → Na^+ + Cl^-

  • P_{solution} = (0.70)*(17.5 \text{ torr}) = 12.3 \text{ torr} (RIGHT)

Raoult’s Law in Action

• Vapor pressure of two volatile solvents:

  • A homogeneous distribution.

  • Each will contribute its own vapor pressure based on its mol fraction (X{sol} or \chi{sol}).

• Let’s look at a 5% w/v mixture of acetic acid:

  • P_{\text{acetic acid}} = 13 \text{ torr at } 20 ˚C

  • P{H2O} = 17 \text{ torr at } 20 ˚C

Boiling Point and Freezing Point Changes

• To find the change in boiling point or freezing point, use the van’t Hoff factor (i).

• Experimentally determined constant for each solvent.

Actual van’t Hoff Factor

• Not all electrolytes fully dissociate.

• The actual van’t Hoff factor is a ratio of measured over theoretical temperature changes for each specific solute.

• Concentration dependent.

Osmosis

• Semi-permeable membranes:

  • Allow for transport of solutes or solvents to pass based on size.

  • Solvents always move from low to high solute concentration.

Osmosis

• Osmotic pressure (Π):

  • The pressure at which osmosis stops.

  • \Pi V = inRT

  • Recall, n/V is molarity.

Colloids

• Colloids – suspension of particles larger than ions, but too small to settle by gravity.

• Particles range from 5-1000 nm.

Colloids

• Colloids – suspension of particles larger than ions, but too small to settle by gravity.

• Colloids scatter light while solutions will allow for transmission or absorption; this is known as the Tyndall effect.

Colloids

• Some compounds (lipids) have hydrophobic and hydrophilic ends.

• This allows for emulsification of fats and oils in an aqueous solution.

• Emulsifier – assists in dissolving a solute in a solvent it doesn’t normally.

Brownian Motion

• Molecules travel in a straight line until they collide.

• Molecules travel with an average speed inversely proportional to their mass.

• Mean free path – average distance molecules travel between collisions.

• Mean free path is smaller for larger molecules.