General Chemistry Practice Final Exam Study Notes

Liquids, Solids, and Intermolecular Forces

Vapor Pressure and Boiling Point

Vapor pressure is the pressure exerted by a vapor in equilibrium with its liquid phase at a given temperature. The boiling point of a liquid is the temperature at which its vapor pressure equals atmospheric pressure.

• Vapor Pressure: Increases with temperature as more molecules have enough energy to escape the liquid phase.

• Boiling Point: The temperature at which vapor pressure equals 1 atm (normal boiling point).

• Intermolecular Forces: Liquids with strong intermolecular forces (e.g., hydrogen bonding) have lower vapor pressures and higher boiling points.

• Graph Interpretation: The provided graph compares vapor pressure curves for water, ethylene glycol, ethanol, and diethyl ether, showing their boiling points and relative vapor pressures at different temperatures.

• Example: Water has a higher boiling point than ethanol or diethyl ether due to stronger hydrogen bonding.

Substance	Normal Boiling Point (°C)	Intermolecular Forces
Diethyl ether	34.6	London dispersion, dipole-dipole
Ethanol	78.3	Hydrogen bonding
Water	100	Hydrogen bonding
Ethylene glycol	197	Hydrogen bonding

Additional info: The graph demonstrates how substances with stronger intermolecular forces require higher temperatures to reach the same vapor pressure as those with weaker forces.

Thermochemistry

Enthalpy Calculations

Enthalpy ($\Delta H$) is a measure of the heat content of a system at constant pressure. It is used to quantify the heat absorbed or released during chemical reactions.

• Standard Enthalpy Change ($\Delta H^\circ$): The enthalpy change measured under standard conditions (1 atm, 25°C).

• Calculation: $\Delta H_{rxn} = \sum \Delta H_f^\circ (\text{products}) - \sum \Delta H_f^\circ (\text{reactants})$

• Units: kJ/mol

• Example: If the enthalpy of formation of reactants and products are given, use the above formula to calculate the reaction enthalpy.

Solutions and Aqueous Reactions

Concentration Units

Concentration describes the amount of solute dissolved in a given quantity of solvent or solution.

• Molarity (M): $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Other Units: Molality, percent by mass, and mole fraction are also used in specific contexts.

• Example: To prepare a 1 M NaCl solution, dissolve 1 mole of NaCl in enough water to make 1 liter of solution.

Chemical Kinetics

Rate Laws and Reaction Order

Chemical kinetics studies the speed of chemical reactions and the factors affecting them.

• Rate Law: $\text{Rate} = k[A]^m[B]^n$

• Order of Reaction: The sum of the exponents in the rate law ($m + n$).

• Units of Rate Constant (k): Depend on the overall order of the reaction.

• Example: For a first-order reaction, $\text{Rate} = k[A]$.

Acids and Bases

pH and pOH Calculations

The pH scale measures the acidity or basicity of a solution. It is related to the concentration of hydrogen ions ($[H^+]$).

• pH: $pH = -\log[H^+]$

• pOH: $pOH = -\log[OH^-]$

• Relationship: $pH + pOH = 14$ (at 25°C)

• Example: If $[H^+] = 1 \times 10^{-3}$ M, then $pH = 3$.

Free Energy and Thermodynamics

Gibbs Free Energy

Gibbs free energy ($\Delta G$) determines the spontaneity of a process at constant temperature and pressure.

• Formula: $\Delta G = \Delta H - T\Delta S$

• Spontaneity: If $\Delta G < 0$, the process is spontaneous; if $\Delta G > 0$, it is nonspontaneous.

• Example: Calculate $\Delta G$ using enthalpy and entropy values at a given temperature.

Additional info:

Due to the limited legibility and context of the original file, some topics and examples have been inferred based on standard General Chemistry curriculum and the visible graph. The file appears to be a practice final exam with answer key, covering a range of topics including phase changes, thermochemistry, solutions, kinetics, acids and bases, and thermodynamics.