solutions
Objectives of the Unit
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.
Mixtures and Solutions
Pure substances are rare; most materials are mixtures of two or more pure substances.
Properties depend on their composition:
Example: Properties of brass (Cu and Zn) differ from those of bronze (Cu and Sn).
Fluoride ions in water can prevent tooth decay (1 ppm) but can be toxic at higher concentrations (1.5 ppm).
Intravenous injections require specific ionic concentrations to match blood plasma.
Focus on liquid solutions and their properties in this unit:
Types of solutions (gaseous, liquid, solid).
Properties like vapor pressure and colligative properties.
Types of Solutions
General Features
Homogeneous mixtures of two or more components.
Composition and properties are uniform.
Largest quantity component is called the solvent.
Remaining components are solutes.
Binary solutions (only two components) are focused here.
Types of Solutions by State
Gaseous Solutions
Gas in gas (e.g., mixture of O2 and N2).
Liquid in gas (e.g., chloroform in nitrogen).
Solid in gas (e.g., camphor in nitrogen).
Liquid Solutions
Gas in liquid (e.g., oxygen in water).
Liquid in liquid (e.g., ethanol in water).
Solid in liquid (e.g., glucose in water).
Solid Solutions
Gas in solid (e.g., hydrogen in palladium).
Liquid in solid (e.g., amalgam of mercury with sodium).
Solid in solid (e.g., copper in gold).
Expressing Concentration of Solutions
Mass Percentage (w/w)
Defined as: [\text{Mass % of a component} = \frac{\text{Mass of the component in solution}}{\text{Total mass of the solution}} \times 100]
Example: A 10% glucose solution means 10g glucose in 90g water, totaling 100g.
Volume Percentage (V/V)
Defined as: [\text{Volume % of a component} = \frac{\text{Volume of the component}}{\text{Total volume of solution}} \times 100]
Example: 10% ethanol in water means 10 mL ethanol in a total solution volume of 100 mL.
Mass by Volume Percentage (w/V)
Mass of solute in 100 mL of solution.
Parts per Million (ppm)
Defined for trace quantities: [\text{ppm} = \frac{\text{Number of parts of the component}}{\text{Total number of parts of all components}} \times 10^6]
Mole Fraction
Defined as: [x_i = \frac{n_i}{n_{total}}]
Useful in relating physical properties like vapor pressure to concentration.
Example Calculations
Calculating Mole Fraction
For 20% C2H6O2 by mass in a 100 g solution:
20 g of C2H6O2.
80 g H2O.
Molar mass of C2H6O2 = 62 g/mol.
Moles of C2H6O2 = 20 g / 62 g/mol = 0.322 mol.
Moles of H2O = 80 g / 18 g/mol = 4.444 mol.
Mole fraction of C2H6O2 = 0.322 / (0.322 + 4.444) = 0.068.
Molarity (M)
Defined as: [M = \frac{\text{Moles of solute}}{\text{Volume of solution in litres}}]
Example Calculation:
5 g NaOH in 450 mL solution:
Moles of NaOH = 5 g / 40 g/mol = 0.125 mol.
Volume in L = 0.450 L.
Molarity = 0.125 mol / 0.450 L = 0.278 M.
Molality (m)
Defined as: [m = \frac{\text{Moles of solute}}{\text{Mass of solvent in kg}}]
Example Calculation:
For 2.5 g ethanoic acid in 75 g benzene:
Moles of C2H4O2 = 2.5 g / 60 g/mol = 0.0417 mol.
Mass of benzene = 75 g / 1000 = 0.075 kg.
Molality = 0.0417 mol / 0.075 kg = 0.556 mol/kg.
Solubility
Depends on the solute and solvent nature as well as temperature and pressure.
Effect of Temperature on Solubility
Solubility of solids: Typically increases with temperature (endothermic processes).
Solubility of gases: Typically decreases with temperature (exothermic processes).
Effect of Pressure on Solubility
Pressure has limited effect on solubility of solids in liquids.
In contrast, increasing pressure increases the solubility of gases in liquids, as stated by Henry's law: [p = K_H x]
Henry's Law and Applications
Henry's law states solubility of a gas in a liquid is proportional to the partial pressure of the gas.
Applications include:
CO2 solubility in soft drinks (sealed under pressure).
Scuba diving (dissolved gases and pressure).
High altitudes (lower oxygen pressure).
Raoult's Law
States that vapor pressure of each component in a mixture is proportional to its mole fraction in the solution: [p_i = x_i p_i^0]
Distinction between ideal and non-ideal solutions based on their obediences:
Ideal solutions conform throughout concentrations.
Non-ideal solutions show deviations (positive or negative).
Colligative Properties
Include elevation of boiling point, depression of freezing point, and osmotic pressure, all depending on particle concentration.
Practical Examples and Calculations
Depression and Elevation of Boiling Point and Freezing Point
Change in boiling point and freezing point directly proportional to the number of solute particles: [\Delta T_b = K_b m] [\Delta T_f = K_f m]
Osmotic Pressure
Defined as the pressure required to prevent osmosis: [P = C R T]
Used for determining molar mass of solutes, especially in proteins and polymers.
Abnormal Molar Masses
Occur due to dissociation/association of solutes, affecting measured properties:
Van't Hoff factor (i) accounts for this behavior: [i = \frac{\text{Normal molar mass}}{\text{Abnormal molar mass}}]
Summary of Key Terms and Concepts
Solutions exist in various forms (solid, liquid, gas).
Concentrations can be expressed in multiple ways (mass %, molarity, molality, etc.).
Key laws (Henry's, Raoult's) govern the behavior of solutions, while colligative properties depend on particle number.
Understanding molarity and its application is crucial for calculations in chemistry.