General Chemistry Practice Exam Study Guide: Thermochemistry, Bonding, Gases, and More

Thermochemistry and Chemical Energy

Internal Energy and Enthalpy

The internal energy of a system is the sum of the kinetic and potential energies of all its components. Enthalpy (H) is a thermodynamic quantity equivalent to the total heat content of a system at constant pressure.

• Internal Energy (U): $U = KE + PE$ (kinetic + potential energy of all particles)

• Enthalpy Change (ΔH): $\Delta H = \Delta U + P\Delta V$

• Exothermic Process: Releases heat to surroundings; ΔH < 0

• Endothermic Process: Absorbs heat from surroundings; ΔH > 0

• Example: Mixing HCl and NaOH is exothermic; heat is released.

Calculating ΔH and Calorimetry

ΔH for reactions can be calculated using enthalpy of formation, calorimetry, or Hess's Law.

• Standard Enthalpy of Formation (ΔHf°): Enthalpy change when 1 mole of a compound forms from its elements in standard states.

• Calorimetry: Measures heat flow using specific heat capacity ($q = mC\Delta T$)

• Example Calculation: For $C_2H_5OH$ with $C = 2.44 \text{ J/g·K}$, heat required for temperature change: $q = mC\Delta T$

Work and Heat in Chemical Reactions

Work is done by or on the system during volume changes at constant pressure.

• Work (w): $w = -P\Delta V$

• At constant pressure: $q_p = \Delta H$

• At constant volume: $q_v = \Delta U$

Chemical Bonding and Molecular Structure

Types of Chemical Bonds

Chemical bonds include ionic, covalent, and metallic bonds. Covalent bonds involve sharing electrons between atoms.

• Single, Double, Triple Bonds: Maximum number of bonds between two atoms is three (triple bond).

• Bond Polarity: Determined by difference in electronegativity.

• Lewis Structures: Show arrangement of electrons; best structure minimizes formal charges.

• Example: For OCl2, choose the Lewis structure with correct octet and formal charges.

Resonance and Formal Charge

Resonance structures represent delocalized electrons; formal charge helps identify the most stable structure.

• Resonance: Multiple valid Lewis structures for a molecule.

• Formal Charge: $\text{FC} = \text{valence electrons} - \text{nonbonding electrons} - \frac{1}{2}\text{bonding electrons}$

• Example: NCN- has several resonance forms; valid forms distribute charges appropriately.

Hybridization and Molecular Geometry

Atomic orbitals combine to form hybrid orbitals, determining molecular geometry.

• sp, sp2, sp3 Hybridization: Number of electron domains determines hybridization.

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

• Example: CO2 is linear (sp hybridization); CH4 is tetrahedral (sp3).

Gases: Properties and Behavior

Gas Laws

Gas behavior is described by several laws relating pressure, volume, temperature, and amount.

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

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

• Ideal Gas Law: $PV = nRT$

• Example: Calculate volume of 1.0 mol gas at STP: $V = \frac{nRT}{P}$

Gas Mixtures and Partial Pressures

In mixtures, each gas exerts a partial pressure; total pressure is the sum of partial pressures.

• Dalton's Law: $P_{\text{total}} = P_1 + P_2 + ... + P_n$

• Collecting Gases Over Water: Account for vapor pressure of water.

Intermolecular Forces and Phase Changes

Types of Intermolecular Forces

Intermolecular forces affect boiling points, melting points, and solubility.

• London Dispersion Forces: Present in all molecules, especially nonpolar.

• Dipole-Dipole Interactions: Between polar molecules.

• Hydrogen Bonding: Strongest, occurs when H is bonded to N, O, or F.

• Example: H2O has hydrogen bonding; CH4 has only London forces.

Phase Changes and Calorimetry

Phase changes involve energy transfer; calorimetry measures heat during these changes.

• Specific Heat Capacity: Amount of heat required to raise temperature of 1 g by 1°C.

• Example: Calculate heat required to raise temperature of a metal sample.

Chemical Kinetics and Equilibrium

Reaction Rates and Mechanisms

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

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

• Activation Energy: Minimum energy required for reaction.

Chemical Equilibrium

At equilibrium, forward and reverse reaction rates are equal.

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

• Le Chatelier's Principle: System shifts to counteract changes in concentration, pressure, or temperature.

Atomic Structure and Periodicity

Electronic Structure of Atoms

Atoms consist of protons, neutrons, and electrons; electron configuration determines chemical properties.

• Quantum Numbers: Describe electron energy levels and orbitals.

• Periodic Trends: Atomic radius, ionization energy, and electronegativity vary across the periodic table.

Tables

Sample Table: Standard Enthalpy of Formation Values

Compound	ΔHf° (kJ/mol)
CO2(g)	-393.5
H2O(l)	-285.8
CH4(g)	-74.8
O2(g)	0

Additional info: Table values inferred from standard enthalpy tables commonly used in general chemistry.

Additional Info

Additional info: These study notes cover topics from thermochemistry, chemical bonding, gas laws, intermolecular forces, and atomic structure, as reflected in the practice exam questions. All equations are provided in LaTeX format for clarity.