solutions

Learning Objectives

  • 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.

Introduction to Solutions

  • Definition of Solutions: Solutions are homogeneous mixtures of two or more components; their utility depends on their composition.

  • Importance of Composition: For instance, the properties of brass differ from those of German silver or bronze due to variations in composition.

  • Concentration of Solutions: Concentrations can affect biological and chemical processes, e.g., fluoride concentrations in water related to dental health.

  • Types of Solutions: Focus is primarily on liquid solutions and their properties.

1.1 Types of Solutions

  • Homogeneous Mixtures: All components are uniformly distributed.

    • Solvent: The component present in the largest quantity, determining the physical state of the solution.

    • Solutes: Components other than the solvent (may be solid, liquid, or gas) that are dissolved.

  • Binary Solutions: Consideration of solutions consisting of just two components (solvent and solute).

Types of Solutions

Type of Solution

Solute Type

Solvent Type

Common Examples

Gaseous Solutions

Gas

Gas

Mixture of oxygen and nitrogen gases

Liquid Solutions

Gas

Liquid

Oxygen dissolved in water

Liquid Solutions

Liquid

Liquid

Ethanol dissolved in water

Solid Solutions

Solid

Liquid

Glucose dissolved in water

Solid Solutions

Gas

Solid

Solution of hydrogen in palladium

Solid Solutions

Solid

Solid

Amalgam of mercury with sodium

Solid Solutions

Solid

Solid

Copper dissolved in gold

Concentration of Solutions

  • Concentration can be described qualitatively (dilute or concentrated) or quantitatively using various units:

    • Mass Percentage (w/w):

      • Given by the formula:Mass % = (Mass of component / Total mass of solution) × 100

      Example: 10% glucose solution = 10 g glucose in 90 g water.

    • Volume Percentage (V/V):

      • Volume % = (Volume of component / Total volume of solution) × 100.

    • Mass by Volume Percentage (w/V):

      • Common in medicine; mass per 100 mL.

    • Parts Per Million (ppm):

      • Ppm = (Number of parts of component / Total number of parts of solution) × 1,000,000.

    • Mole Fraction (x):

      • Mole fraction of a component = (Number of moles of component / Total number of moles of all components).

    • Molarity (M):

      • Defined as moles of solute per litre of solution: [ M = \frac{Moles \ of \ solute}{Volume \ of \ solution \ (litres)} ]

    • Molality (m):

      • Moles of solute per kilogram of solvent.

Calculation of Concentrations

  1. Mole Fraction Calculation Example: 20% mass ethylene glycol in 100 g solution:

    • Moles of C2H6O2 = 20 g / 62 g/mol = 0.322 mol.

    • Moles of water = 80 g / 18 g/mol = 4.444 mol.

    • Mole Fraction of C2H6O2: ( x = \frac{0.322}{0.322 + 4.444} = 0.068 )

    • Mole Fraction of water: ( 1 - x = 0.932 )

Molarity Example

  • Calculating Molarity of NaOH: 5 g in 450 mL:

    • Moles of NaOH = 5 g / 40 g/mol = 0.125 mol.

    • Molarity ( M = \frac{0.125 \ mol}{0.450 \ litre} = 0.278 \ mol/L )

Solubility Factors

  • Solubility depends on: nature of solute/solvent, temperature, and pressure.

  • Dissolution Process: Solute dissolves, leading to an equilibrium between solute molecules and undissolved solute.

Temperature Effect on Solubility

  • Generally, solubility of solids in liquids increases with temperature; for gases, it decreases.

  • Henry’s Law: Relates gas solubility to pressure:

    • ( P = K_H \cdot x )

    • P = partial pressure of gas, K_H = Henry's law constant, x = mole fraction in solution.

Applications of Henry’s Law

  1. Soft Drinks: CO2 is solubilized under pressure.

  2. Scuba Diving: Increased gas solubility at depth leads to decompression sickness when returning to surface.

  3. High Altitude: Reduced oxygen availability leads to symptoms like anoxia.

Raoult’s Law

  • States that the partial vapor pressure of each component in a solution is proportional to its mole fraction:

    • For volatile components:

    • ( P_{total} = P_1^0 x_1 + P_2^0 x_2 )

  • Ideal Solutions: Where components behave according to Raoult’s law.

  • Deviations: Positive (higher than Raoult's prediction) and negative (lower than Raoult's prediction).

Colligative Properties

  • Properties that depend on the number of solute particles, not their identity:

    1. Lowering of vapor pressure.

    2. Elevation of boiling point.

    3. Depression of freezing point.

    4. Osmotic pressure.

    • All described mathematically:

      • E.g., ( \Delta T_b = i K_b m )

Osmotic Pressure and Reverse Osmosis

  • Defined as the pressure required to prevent osmosis; connected to solute concentration.

  • Reverse osmosis is used for water purification, e.g., desalination.

Abnormal Colligative Properties

  • Occurs when solutes dissociate or associate, impacting molar mass calculations.

  • Van't Hoff Factor (i):

    • Explains discrepancies.

Summary of Key Concepts

  • A solution is a homogeneous mixture of substances.

  • Concentrations can be expressed in various terms: mole fraction, molarity, molality, and percentage.

  • Colligative properties are essential in understanding solution behaviors.

  • Henry’s law and Raoult’s law describe gas solubility and vapor pressure trends.