General Chemistry: Solutions, Intermolecular Forces, and Colligative Properties – Study Notes

General Chemistry: Solutions, Intermolecular Forces, and Colligative Properties

Exam Structure and Policies

This exam consists of multiple-choice and free-response sections, covering key topics in general chemistry. Students are expected to follow academic integrity policies and use only approved materials during the exam.

• Multiple Choice Section: Each question is worth 3 points. Students may skip up to three questions without penalty. No partial credit is given.

• Free Response Section: Worth 20 points in total. Partial credit is awarded based on shown work.

• Allowed Materials: Scratch paper, equation sheet, and periodic table provided. No programmable calculators or electronic devices allowed.

• Exam Duration: 90 minutes. Only the free response sheet is submitted at the end.

Intermolecular Forces

Types and Strengths of Intermolecular Forces

Intermolecular forces are the attractions between molecules, influencing physical properties such as boiling and melting points, solubility, and vapor pressure.

• London Dispersion Forces (LDF): Present in all molecules, especially significant in nonpolar molecules. Caused by temporary dipoles.

• Dipole-Dipole Forces: Occur between polar molecules with permanent dipoles.

• Hydrogen Bonding: A strong type of dipole-dipole interaction, occurs when H is bonded to N, O, or F.

• Ion-Dipole Forces: Occur between ions and polar molecules, important in solutions of ionic compounds in polar solvents.

Example: SO2 exhibits strong dipole-dipole forces due to its bent shape and polar bonds.

Identifying Intermolecular Forces in Molecules

• Nonpolar molecules (e.g., Cl2, Br2): Only LDF present.

• Polar molecules (e.g., SO2): Dipole-dipole and LDF present.

• Molecules with N-H, O-H, or F-H bonds (e.g., CH3CH2NH2): Hydrogen bonding, dipole-dipole, and LDF present.

States of Matter

Properties of Solids, Liquids, and Gases

The three main states of matter differ in shape and volume characteristics:

• Solid: Definite shape and volume.

• Liquid: Definite volume, indefinite shape.

• Gas: Indefinite shape and volume.

Phase Changes

• Deposition: The process where a gas becomes a solid directly, skipping the liquid phase.

• Other phase changes: Sublimation (solid to gas), melting (solid to liquid), condensation (gas to liquid), freezing (liquid to solid).

Solutions and Solubility

Solubility and Saturation

Solubility is the maximum amount of solute that can dissolve in a solvent at a given temperature.

• Unsaturated Solution: Contains less solute than the maximum amount that can dissolve.

• Saturated Solution: Contains the maximum amount of dissolved solute; additional solute will not dissolve.

• Supersaturated Solution: Contains more solute than is normally possible at that temperature; unstable.

Example: At 40°C, 100 g of water can dissolve about 62 g of KNO3. A solution with 20 g KNO3 is unsaturated.

Factors Affecting Solubility

• Temperature: Solubility of solids generally increases with temperature; for gases, solubility decreases with increasing temperature.

• Pressure: Solubility of gases increases with pressure (Henry's Law).

• "Like Dissolves Like": Polar solvents dissolve polar/ionic solutes; nonpolar solvents dissolve nonpolar solutes.

Solubility Curves

Solubility curves show how the solubility of a substance changes with temperature. They are used to determine if a solution is saturated, unsaturated, or supersaturated at a given temperature.

Concentration Units

Molarity (M)

Molarity is the number of moles of solute per liter of solution.

• Formula: $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Example: Dissolving 200 g NH4Cl (molar mass = 53.50 g/mol) in 1 L solution: $M = \frac{200}{53.50} = 3.74$ M

Mass Percent

Mass percent is the mass of solute divided by the total mass of solution, multiplied by 100%.

• Formula: $\text{Mass percent} = \frac{\text{mass of solute}}{\text{mass of solution}} \times 100\%$

• Example: 47.3 g AgI in 473 mL solution (density = 1.12 g/mL): $\text{Mass percent} = \frac{47.3}{473 \times 1.12} \times 100\% = 9.03\%$

Mole Fraction

Mole fraction is the ratio of moles of a component to the total moles in the solution.

• Formula: $X_A = \frac{\text{moles of A}}{\text{total moles}}$

• Example: 14.6 g ethanol (C2H5OH, 46.0 g/mol) and 53.6 g H2O (18.0 g/mol): $X_{\text{ethanol}} = \frac{0.317}{2.97 + 0.317} = 0.0962$

Colligative Properties

Boiling Point Elevation

Adding a nonvolatile solute to a solvent raises its boiling point.

• Formula: $\Delta T_b = i K_b m$

• Where: $i$ = van't Hoff factor, $K_b$ = boiling point elevation constant, $m$ = molality

• Example: For 1.0 m solution, $\Delta T_b = (0.512 \frac{\text{°C}}{\text{m}})(1.0 \text{ m}) = 0.512$°C

Osmotic Pressure

Osmotic pressure is the pressure required to stop the flow of solvent into a solution through a semipermeable membrane.

• Formula: $\Pi = iMRT$

• Where: $\Pi$ = osmotic pressure, $i$ = van't Hoff factor, $M$ = molarity, $R$ = gas constant, $T$ = temperature in Kelvin

• Example: For a protein solution: $M = \frac{\Pi}{RT}$, then use mass and volume to find molar mass.

Vapor Pressure Lowering (Raoult's Law)

The vapor pressure of a solution is lower than that of the pure solvent.

• Formula: $P = X_{\text{solvent}} P^0_{\text{solvent}}$

• Where: $P$ = vapor pressure of solution, $X_{\text{solvent}}$ = mole fraction of solvent, $P^0_{\text{solvent}}$ = vapor pressure of pure solvent

• Example: $P = 0.792 \times 23.8 = 18.8$ mm Hg

Thermochemistry of Solutions

Heat of Solution

The process of dissolving can be exothermic or endothermic, depending on the relative strengths of solute-solute, solvent-solvent, and solute-solvent interactions.

• Exothermic: Heat is released; solution warms up.

• Endothermic: Heat is absorbed; solution cools down.

• Example: If 20.8 kJ is released when 200 g of solute is dissolved, the process is exothermic.

Gas Solubility and Henry's Law

Henry's Law

The solubility of a gas in a liquid is directly proportional to the partial pressure of the gas above the liquid.

• Formula: $S = k_H P$

• Where: $S$ = solubility, $k_H$ = Henry's law constant, $P$ = partial pressure of gas

• Example: For CO2 in soda at 25°C, $k_H = 3.4 \times 10^{-2} \frac{\text{mol}}{\text{L} \cdot \text{atm}}$, $P = 7.0$ atm: $S = (3.4 \times 10^{-2})(7.0) = 0.24$ M

Homogeneous Solutions

Formation of Homogeneous Solutions

Homogeneous solutions form when solute and solvent are compatible in terms of polarity and intermolecular forces.

• "Like Dissolves Like": Polar solutes dissolve in polar solvents; nonpolar in nonpolar.

• Example: NaCl (ionic, polar) dissolves in CH3OH (polar); C6H6 (nonpolar) dissolves in CCl4 (nonpolar).

Summary Table: Types of Intermolecular Forces

Type of Force	Occurs Between	Relative Strength	Example
London Dispersion (LDF)	All molecules	Weakest	Cl2, Br2
Dipole-Dipole	Polar molecules	Intermediate	SO2
Hydrogen Bonding	H bonded to N, O, or F	Strong	H2O, CH3CH2NH2
Ion-Dipole	Ions and polar molecules	Strongest (in solution)	Na+ in H2O

Key Equations and Constants

• Molarity: $M = \frac{\text{mol solute}}{\text{L solution}}$

• Mass Percent: $\text{Mass percent} = \frac{\text{mass solute}}{\text{mass solution}} \times 100\%$

• Mole Fraction: $X_A = \frac{\text{mol A}}{\text{total mol}}$

• Boiling Point Elevation: $\Delta T_b = i K_b m$

• Osmotic Pressure: $\Pi = iMRT$

• Raoult's Law: $P = X_{\text{solvent}} P^0_{\text{solvent}}$

• Henry's Law: $S = k_H P$

Additional info: These notes synthesize the main concepts, equations, and applications from the provided exam questions and solutions, offering a comprehensive review for students preparing for a general chemistry exam on solutions, intermolecular forces, and colligative properties.