General Chemistry Practice Exam Study Guide

Thermochemistry and Chemical Energy

Internal Energy, Enthalpy, and Calorimetry

Thermochemistry studies the energy changes that occur during chemical reactions and physical changes. Key concepts include internal energy, enthalpy, and calorimetry.

• Internal Energy (U): The total energy contained within a system, including both kinetic and potential energies of all particles.

• Enthalpy (H): The heat content of a system at constant pressure. The change in enthalpy ($\Delta H$) is used to quantify heat flow in chemical reactions at constant pressure.

• Calorimetry: The measurement of heat flow. A calorimeter is used to determine the heat exchanged in a reaction.

• Specific Heat Capacity (c): The amount of heat required to raise the temperature of 1 gram of a substance by 1°C.

Key Equations:

• Heat transfer: $q = mc\Delta T$

• Work at constant pressure: $w = -P\Delta V$

• Relationship between internal energy, heat, and work: $\Delta U = q + w$

• Enthalpy change: $\Delta H = \Delta U + P\Delta V$

Example: Calculating the heat absorbed by a metal sample using its mass, specific heat, and temperature change.

Chemical Reactions and Stoichiometry

Types of Reactions and Energy Changes

Chemical reactions can be classified by the energy changes involved and the types of substances reacting.

• Endothermic Process: Absorbs heat from the surroundings ($\Delta H > 0$).

• Exothermic Process: Releases heat to the surroundings ($\Delta H < 0$).

• Combustion Reactions: Rapid reactions with oxygen that release energy.

• Standard Enthalpy of Formation ($\Delta H_f^\circ$): The enthalpy change when one mole of a compound is formed from its elements in their standard states.

Example: Calculating $\Delta H$ for a reaction using standard enthalpies of formation:

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

Atomic Structure and Periodicity

Electron Configuration and Periodic Trends

The arrangement of electrons in atoms determines chemical properties and periodic trends.

• Electron Configuration: The distribution of electrons among atomic orbitals.

• Periodic Trends: Properties such as atomic radius, ionization energy, and electronegativity show predictable patterns across periods and groups.

Example: Determining the electron configuration of an element and predicting its chemical behavior.

Chemical Bonding and Molecular Structure

Ionic and Covalent Bonding

Chemical bonds form when atoms share or transfer electrons to achieve stable electron configurations.

• Ionic Bonds: Formed by the transfer of electrons from metals to nonmetals, resulting in oppositely charged ions.

• Covalent Bonds: Formed by the sharing of electrons between nonmetal atoms.

• Lewis Structures: Diagrams that show the arrangement of valence electrons in molecules.

• Resonance: Some molecules can be represented by two or more valid Lewis structures.

Example: Drawing the Lewis structure for COCl2 and identifying resonance forms for NCO-.

Bonding Theories and Molecular Geometry

• VSEPR Theory: Predicts the shapes of molecules based on electron pair repulsion.

• Hybridization: Atomic orbitals mix to form new hybrid orbitals (e.g., sp, sp2, sp3).

• Bond Order: The number of chemical bonds between a pair of atoms.

Example: Determining the geometry of SF6 (octahedral) and the hybridization of carbon in CH2O (sp2).

Gases: Properties and Behavior

Gas Laws and Calculations

Gases exhibit predictable behavior described by several laws.

• Ideal Gas Law: $PV = nRT$

• Boyle's Law: $P_1V_1 = P_2V_2$ (at constant T and n)

• Charles's Law: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$ (at constant P and n)

• Avogadro's Law: $V \propto n$ (at constant T and P)

• Dalton's Law of Partial Pressures: $P_{total} = \sum P_i$

Example: Calculating the volume of a gas at STP or the partial pressure of a component in a mixture.

Intermolecular Forces and States of Matter

Liquids, Solids, and Phase Changes

Intermolecular forces determine the physical properties of substances.

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

• Boiling and Melting Points: Influenced by the strength of intermolecular forces.

• Phase Changes: Transitions between solid, liquid, and gas phases involve energy changes.

Example: Comparing boiling points of different substances based on their intermolecular forces.

Solutions and Their Properties

Concentration and Colligative Properties

Solutions are homogeneous mixtures with properties dependent on the concentration of solute particles.

• Concentration Units: Molarity (M), molality (m), percent composition.

• Colligative Properties: Properties that depend on the number of solute particles (e.g., boiling point elevation, freezing point depression).

Example: Calculating the molarity of a solution or the boiling point elevation caused by a solute.

Chemical Kinetics and Equilibrium

Reaction Rates and Dynamic Equilibrium

Chemical kinetics studies the speed of reactions, while equilibrium describes the balance between forward and reverse reactions.

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

• Equilibrium Constant (K): $K = \frac{[\text{products}]}{[\text{reactants}]}$ (at equilibrium)

Example: Determining the effect of concentration changes on the position of equilibrium.

Tables

Sample Table: Comparison of Intermolecular Forces

Type of Force	Relative Strength	Example
London Dispersion	Weakest	He, Ar
Dipole-Dipole	Intermediate	HCl, SO2
Hydrogen Bonding	Strongest (of IMFs)	H2O, NH3

Sample Table: Hybridization and Geometry

Hybridization	Geometry	Bond Angle
sp	Linear	180°
sp2	Trigonal planar	120°
sp3	Tetrahedral	109.5°

Additional info:

• This study guide covers topics from thermochemistry, atomic structure, chemical bonding, molecular geometry, gases, intermolecular forces, solutions, and chemical equilibrium, as reflected in the exam questions.

• Sample calculations, definitions, and tables have been added for academic completeness.