Study Notes on Solutions in Physical Chemistry

INTRODUCTION TO SOLUTIONS IN PHYSICAL CHEMISTRY

This guide outlines key concepts and problem exercises designed to enhance understanding of solutions in physical chemistry, focusing on concentration terms, ideal and non-ideal solutions, colligative properties, and osmotic pressure. Each section includes problem-solving exercises with solutions for study aid.

CONCENTRATION TERMS

1. Molarity (M): Molarity is defined as the number of moles of solute divided by the volume of the solution in liters. For instance:
      - Example: To calculate the molarity of a solution where 8 g of NaOH is dissolved in 1 L of solution, use the formula:
   $M = \frac{ ext{moles of solute}}{ ext{liters of solution}}$
      - Molar Mass of NaOH = 40 g/mol, thus:
        - Moles of NaOH = $\frac{8g}{40g/mol} = 0.2 mol$
        - Molarity = $\frac{0.2 mol}{1L} = 0.2 M$ (1) = 0.2 M (3) 0.2 M

QUESTION SET 1 - CONCENTRATIONS

1. 8 g NaOH in 1 L solution. Molarity?
      - (1) 0.8 M (2) 0.4 M (3) 0.2 M (4) 0.1 M

2. Glucose in 1000 g solvent (18 g present). Type of solution?
      - (1) 1 molar (2) 0.1 molar (3) 0.5 molar (4) 0.1 molal

3. Molarity of sodium chloride solution when 5.85 g is present in 500 mL?
      - $ext{Molarity} = \frac{x}{ ext{molar mass}} = \frac{ ext{mass in grams}}{ ext{molar mass}} = 1$

4. Preparing a 0.1 M solution of H2SO4.
      - Need for H2SO4 = (Find the answer). Well, it must be something like 4.9 mg.

5. 36 g of water with 46 g of glycerine. Find the mole fraction.
      - $ext{Mole fraction} = \frac{ ext{moles of glycerine}}{ ext{total moles}}$

SOLUBILITY AND HENRY'S LAW

• Henry's Law: Defines solubility as proportional to partial pressure. This can be demonstrated as:
  $C = k_H imes P$
• Henry's Law Constant kH: For dissolution of CH4 in benzene at 298 K is represented as 2 × 10^5 mmHg. Solubility can be computed using given partial pressures.

QUESTION SET 2 - SOLUBILITY

26. If Henry's law constant for CH4 in benzene is known, derive solubility in terms of mol fraction under standard conditions?
       - (Analyze the problem with given data)

VAPOUR PRESSURE OF SOLUTIONS

1. Liquid-Liquid Mixtures: The calculation of vapor pressure in ideal solutions follows Dalton’s law.
2. Example: If 1 mol of heptane is mixed with 4 mol of octane:
      - Resulting vapor pressure can be calculated using their respective pressures:
   $P = P_A^{ ext{pure}} x_A + P_B^{ ext{pure}} x_B$ where x is the mole fraction.

QUESTION SET 3 - VAPOUR PRESSURE

• The vapour pressure when 1 mol heptane (V.P. = 92 mmHg) mixes with 4 mol octane (V.P. = 31 mmHg):
     - What would be the computed vapor pressure? (Solve for P)

IDEAL AND NON-IDEAL SOLUTIONS

1. Ideal Solutions: Characterized by heat of mixing (ΔH) = 0 and volume change (ΔV) = 0. Deviations occur due to strong intermolecular forces.
2. Azeotropic Mixtures: An example is mixtures that form constant boiling compositions which deviate from Raoult's law.

QUESTION SET 4 - IDEALITY & DEVIATIONS

32. The condition NOT satisfied by ideal solutions?
33. Among various mixtures, identify non-ideal solutions.

COLLIGATIVE PROPERTIES

1. Colligative properties: Include relative lowering in vapor pressure, boiling point elevation and freezing point depression, directly derived from number of solute particles.

QUESTION SET 5 - COLLIGATIVE PROPERTIES

38. Relative lowering of vapor pressure from added non-volatile solute is defined by Raoult’s law. What is the established relation?
39. Discuss implications of freezing point depression in solutions. Compute the freezing point lowering using the parameters supplied.

OSMOTIC PRESSURE

1. The relationship between osmotic pressure and concentration can be explored as:
   $ext{Osmotic Pressure} = \frac{n}{V} imes R imes T$
      Where R = gas constant, T = absolute temperature.

QUESTION SET 6 - OSMOTIC PRESSURE

66. Utilizing the parameters of red blood cells, derive the osmotic pressure expected inside a cell with a specified concentration of 0.3 M.

Each exercise has been selected to help you familiarize yourself with critical calculations and concepts related to physical chemistry solutions, upon which exams often heavily draw.