General Chemistry Final Exam Review: Structure, Properties, Organic Chemistry, Thermochemistry, Equilibrium, and Electrochemistry

Unit 1: Structure and Properties

Lewis Structures, VSEPR, and Molecular Properties

This section reviews the basics of molecular structure, including Lewis structures, VSEPR notation, molecular shapes, and polarity.

• Lewis Structure: A diagram showing the arrangement of valence electrons among atoms in a molecule.

• VSEPR Notation: Valence Shell Electron Pair Repulsion theory predicts molecular geometry based on electron pair repulsion.

• Molecular Shape: Determined by the number of bonding and lone pairs around the central atom (e.g., linear, trigonal planar, tetrahedral).

• Polarity: A molecule is polar if it has an uneven distribution of electron density, resulting in a dipole moment.

Example Table:

Molecule	Lewis Structure	VSEPR Notation	Molecular Shape	Polar/Nonpolar
H2O	H–O–H (with 2 lone pairs on O)	AX2E2	Bent	Polar
CO2	O=C=O	AX2	Linear	Nonpolar
NH3	H–N–H (with 1 lone pair on N)	AX3E	Trigonal pyramidal	Polar

Additional info: The VSEPR notation AXnEm describes the number of atoms (X) and lone pairs (E) around the central atom (A).

Valence Electron Configuration

• Valence electrons are the outermost electrons involved in bonding.

• Electron configuration determines chemical reactivity and periodic trends.

• Example: Sodium (Na): 1s22s22p63s1

Periodic Trends

• Atomic radius: Increases down a group, decreases across a period.

• Ionization energy: Energy required to remove an electron; increases across a period, decreases down a group.

• Electronegativity: Tendency to attract electrons; increases across a period, decreases down a group.

Unit 2: Organic Chemistry

Naming and Identifying Organic Compounds

This section covers the basics of organic nomenclature, functional groups, and boiling point trends.

• Alkanes: Saturated hydrocarbons with single bonds (e.g., CnH2n+2).

• Alkenes: Unsaturated hydrocarbons with at least one double bond (e.g., CnH2n).

• Alkynes: Unsaturated hydrocarbons with at least one triple bond (e.g., CnH2n-2).

• Functional Groups: Specific groups of atoms within molecules that determine chemical reactivity (e.g., alcohol, ketone, carboxylic acid).

Example Table: Functional Groups

Group	Structure	Suffix/Prefix
Alcohol	–OH	-ol
Aldehyde	–CHO	-al
Ketone	–CO–	-one
Carboxylic Acid	–COOH	-oic acid

• Boiling Point Trends: Influenced by molecular mass and intermolecular forces (hydrogen bonding > dipole-dipole > London dispersion).

Unit 3: Thermochemistry & Rate of Reactions

Thermochemical Equations and Enthalpy

Thermochemistry studies energy changes in chemical reactions, focusing on enthalpy ($\Delta H$).

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

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

Example Equation:

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

Reaction Mechanisms and Rate Laws

• Reaction Mechanism: Sequence of elementary steps by which a reaction occurs.

• Rate Law: Expresses the rate as a function of reactant concentrations: $\text{Rate} = k[A]^m[B]^n$

• Catalysts: Substances that increase reaction rate without being consumed.

Unit 4: Chemical Equilibrium, Solubility, and Acids and Bases

Chemical Equilibrium

At equilibrium, the rates of the forward and reverse reactions are equal, and concentrations remain constant.

• Equilibrium Constant ($K$): $K = \frac{[\text{products}]}{[\text{reactants}]}$ (for a balanced equation).

• Le Châtelier's Principle: If a system at equilibrium is disturbed, it will shift to counteract the disturbance.

Acids and Bases

• Brønsted-Lowry Acid: Proton donor.

• Brønsted-Lowry Base: Proton acceptor.

• pH Calculation: $\text{pH} = -\log[\text{H}^+]$

• Buffer Solutions: Resist changes in pH when small amounts of acid or base are added.

Unit 5: Electrochemistry

Redox Reactions and Electrochemical Cells

Electrochemistry studies the transfer of electrons in chemical reactions, including galvanic (voltaic) and electrolytic cells.

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Electrochemical Cell: Device that converts chemical energy into electrical energy (or vice versa).

• Cell Potential ($E^\circ_{cell}$): $E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}$

Example Table: Activity Series

Metal	Tendency to be Oxidized
Li	Very high
Na	High
Cu	Low

• Spontaneity: A positive $E^\circ_{cell}$ indicates a spontaneous reaction.

Additional info: The review also includes balancing redox reactions, identifying oxidizing/reducing agents, and predicting cell reactions using the activity series.