General Chemistry Final Exam Study Guide: Key Topics and Concepts

General Chemistry Final Exam Study Guide

Stoichiometry and Chemical Reactions

Stoichiometry involves the quantitative relationships between reactants and products in chemical reactions. Understanding balanced chemical equations is essential for calculating the amounts of substances consumed and produced.

• Balanced Chemical Equation: Shows the proportions of reactants and products. Example: $C_6H_{12}O_6 (aq) \rightarrow 2 C_2H_5OH (aq) + 2 CO_2 (g)$

• Mole Concept: The mole is a counting unit for atoms, molecules, or ions. $\text{Moles} = \frac{\text{Mass (g)}}{\text{Molar Mass (g/mol)}}$

• Percent Yield: Measures the efficiency of a reaction. $\text{Percent Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100\%$

• Example: Calculating the mass of ethanol produced from a given mass of glucose using stoichiometry.

Gases and Gas Laws

The behavior of gases is described by several laws that relate pressure, volume, temperature, and amount of gas.

• Ideal Gas Law: $PV = nRT$

• Partial Pressure: The pressure exerted by each gas in a mixture. $P_{\text{total}} = P_1 + P_2 + ... + P_n$

• Root Mean Square Velocity: Describes the average speed of gas molecules. $u_{rms} = \sqrt{\frac{3RT}{M}}$

• Effusion and Diffusion: Effusion is the escape of gas through a small hole; diffusion is the mixing of gases.

• Example: Calculating the molar mass of a gas from density and pressure data.

Thermochemistry

Thermochemistry studies the energy changes that occur during chemical reactions, especially heat transfer.

• Enthalpy Change ($\Delta H$): The heat change at constant pressure.

• Standard Enthalpy of Formation: The enthalpy change when one mole of a compound is formed from its elements in their standard states.

• Hess's Law: The total enthalpy change for a reaction is the sum of the enthalpy changes for individual steps.

• Specific Heat Capacity ($c$): The amount of heat required to raise the temperature of 1 g of a substance by 1°C. $q = mc\Delta T$

• Example: Calculating the heat required to raise the temperature of a sample.

Quantum Mechanics and Atomic Structure

Quantum mechanics explains the behavior of electrons in atoms, including their energy levels and orbitals.

• Aufbau Principle: Electrons fill orbitals starting with the lowest energy first.

• Electron Configuration: Describes the arrangement of electrons in an atom. Example: $1s^2 2s^2 2p^6$

• Quantum Numbers: Define the properties of atomic orbitals and electrons: principal ($n$), angular ($l$), magnetic ($m_l$), and spin ($m_s$).

• Energy of a Photon: $E = h\nu = \frac{hc}{\lambda}$

• Example: Calculating the energy associated with a given wavelength of light.

Periodic Properties of the Elements

The periodic table organizes elements by increasing atomic number and recurring chemical properties.

• Trends: Atomic radius, ionization energy, electron affinity, and electronegativity vary predictably across periods and groups.

• Classification: Metals, nonmetals, and metalloids have distinct properties.

• Example: Ranking elements by atomic radius or ionization energy.

Chemical Bonding and Molecular Structure

Chemical bonds form when atoms share or transfer electrons. Molecular structure determines physical and chemical properties.

• Lewis Structures: Diagrams showing the arrangement of electrons in molecules.

• VSEPR Theory: Predicts molecular shapes based on electron pair repulsion.

• Hybridization: Mixing of atomic orbitals to form new hybrid orbitals.

• Molecular Orbital Theory: Describes bonding using molecular orbitals formed from atomic orbitals.

• Example: Drawing Lewis structures and predicting molecular geometry.

Intermolecular Forces and States of Matter

Intermolecular forces determine the physical properties of substances, such as boiling and melting points.

• Types of Forces: London dispersion, dipole-dipole, hydrogen bonding.

• Phase Diagrams: Graphs showing the state of matter at different temperatures and pressures.

• Example: Identifying the strongest intermolecular force in a compound.

Useful Equations and Constants

Several equations and constants are essential for solving problems in general chemistry.

• Gas Law Equations: $PV = nRT$, $P_1V_1 = P_2V_2$, $\frac{P_1}{T_1} = \frac{P_2}{T_2}$

• Thermochemical Equations: $q = mc\Delta T$, $\Delta H = \sum \Delta H_{\text{products}} - \sum \Delta H_{\text{reactants}}$

• Conversion Factors: 1 atm = 101.3 kPa, 1 L = 1000 mL, 1 mol = 6.022 × 1023 particles

• Constants: $R = 0.0821 \text{ L atm mol}^{-1} \text{K}^{-1}$, $h = 6.626 \times 10^{-34} \text{J s}$, $c = 3.00 \times 10^8 \text{m/s}$

Sample Table: Comparison of Specific Heat Capacity and Heat Capacity

Property	Specific Heat Capacity	Heat Capacity
Definition	Heat required to raise temperature of 1 g of substance by 1°C	Heat required to raise temperature of entire sample by 1°C
Units	J/g·°C	J/°C
Depends on	Nature of substance	Mass and nature of substance

Additional info:

• This guide covers topics from chemical reactions and stoichiometry, gas laws, thermochemistry, quantum mechanics, periodic properties, chemical bonding, intermolecular forces, and phase diagrams, as reflected in the exam questions and provided data sheet.

• Equations and constants are taken from the provided data sheet for reference during calculations.