Solutions

• After studying this Unit, you will be able to:

  • Describe the formation of different types of solutions
  • Express concentration of solution in different units
  • State and explain Henry’s law and Raoult’s law
  • Distinguish between ideal and non-ideal solutions
  • Explain deviations of real solutions from Raoult’s law
  • Describe colligative properties of solutions and correlate these with molar masses of the solutes
  • Explain abnormal colligative properties exhibited by some solutes in solutions
• In normal life, mixtures containing two or more pure substances are more common than pure substances
• The properties of mixtures like brass, German silver, and bronze vary based on their composition
• 1 part per million (ppm) of fluoride ions in water prevents tooth decay, while higher concentrations can be harmful
• Intravenous injections are dissolved in water with specific ionic concentrations matching blood plasma concentrations
• Solutions are homogeneous mixtures of two or more components, with the component present in the largest quantity known as the solvent
• Binary solutions consist of two components, which can be solid, liquid, or gaseous
• Concentration of a solution can be described quantitatively using methods like mass percentage, volume percentage, mass by volume percentage, parts per million, and mole fraction
• Mass percentage (w/w) is defined as the mass of a component in the solution divided by the total mass of the solution, multiplied by 100
• Volume percentage (V/V) is defined as the volume of a component in the solution divided by the total volume of the solution, multiplied by 100
• Mass by volume percentage is the mass of solute dissolved in 100 mL of the solution
• Parts per million (ppm) is used to express trace quantities of a solute in a solution
• Mole fraction is the number of moles of a component divided by the total number of moles of all components in the solution
• Molarity (M) is the number of moles of solute dissolved in one litre (or one cubic decimetre) of solution
• Molality (m) is defined as the number of moles of the solute per kilogram (kg) of the solvent
• Molality (m) = Moles of solute / Mass of solvent in kg
• Example: Calculate the molality of 2.5 g of ethanoic acid (CH3COOH) in 75 g of benzene
• Molar mass of C2H4O2: 12 × 2 + 1 × 4 + 16 × 2 = 60 g mol–1
• Moles of C2H4O2 = 2.5 g / 60 g mol–1 = 0.0417 mol
• Mass of benzene in kg = 75 g / 1000 g kg–1 = 75 × 10–3 kg
• Molality of C2H4O2 = 0.0417 mol / 0.075 kg = 0.556 mol kg–1
• Every solid does not dissolve in a given liquid
• Polar solutes dissolve in polar solvents and non-polar solutes in non-polar solvents
• Solute dissolves in a solvent if the intermolecular interactions are similar
• Dissolution is the process where solute particles dissolve in the solvent
• Crystallisation is the process where solute particles separate out of the solution
• Dynamic equilibrium is reached when the rate of dissolution equals the rate of crystallisation
• Saturated solution is a solution where no more solute can be dissolved at the same temperature and pressure
• Henry's Law states that the solubility of a gas in a liquid is directly proportional to the partial pressure of the gas above the liquid
• Solubility of a substance depends on the nature of the substances, temperature, and pressure
• Henry's Law constant (KH) relates the partial pressure of a gas to its solubility in a liquid
• Solubility is the maximum amount of solute that can be dissolved in a given amount of solvent
• Henry's Law states that the solubility of a gas in a liquid is directly proportional to the pressure of the gas above the solution
• The solubility of gases in liquids decreases with an increase in temperature
• Dissolution of gas molecules in a liquid is an exothermic process
• Raoult's Law states that for a solution of volatile liquids, the partial vapor pressure of each component is directly proportional to its mole fraction in the solution
• Dalton's Law of partial pressures states that the total pressure over a solution phase is the sum of the partial pressures of the components of the solution
• The total vapor pressure over a solution can be related to the mole fraction of any one component
• The total vapor pressure over a solution varies linearly with the mole fraction of a component
• The composition of the vapor phase in equilibrium with the solution is determined by the partial pressures of the components