General Chemistry Exam Review: Key Concepts and Practice Problems

Unit 1: Introduction to Chemistry

Atoms, Molecules, and Ions

This section covers the foundational concepts of chemistry, including the identification of atoms, elements, and molecules, as well as the formation and naming of ions and compounds.

• Atoms are the basic units of matter, consisting of protons, neutrons, and electrons.

• Molecules are groups of two or more atoms bonded together.

• Ions are atoms or molecules that have gained or lost electrons, resulting in a net charge.

• Polyatomic ions are ions composed of more than one atom (e.g., sulfate SO42−, nitrate NO3−).

Example: The formula for sodium sulfate is Na2SO4.

Predicting Charges and Writing Formulas

• Metals typically form positive ions (cations), while nonmetals form negative ions (anions).

• To write the formula for an ionic compound, balance the charges so the total positive and negative charges are equal.

Example: Magnesium chloride: Mg2+ and Cl− combine to form MgCl2.

Unit 1 (Part 1): Atomic Structure and Properties (The Mole)

Atomic Structure

Atoms are composed of protons, neutrons, and electrons. The number of protons defines the element, while isotopes differ in the number of neutrons.

• Atomic number (Z): Number of protons in the nucleus.

• Mass number (A): Total number of protons and neutrons.

• Isotopes: Atoms of the same element with different numbers of neutrons.

Example: Chlorine-35 and Chlorine-37 are isotopes of chlorine.

The Mole and Molar Mass

• The mole (mol) is a counting unit for atoms and molecules. 1 mol = $6.022 \times 10^{23}$ particles (Avogadro's number).

• Molar mass is the mass of one mole of a substance, expressed in grams per mole (g/mol).

Example: The molar mass of CaCl2 is the sum of the atomic masses of Ca and two Cl atoms.

Percent Composition and Empirical Formulas

• Percent composition is the percentage by mass of each element in a compound.

• Empirical formula is the simplest whole-number ratio of atoms in a compound.

Example: A compound with 40% C, 6.7% H, and 53.3% O has an empirical formula of CH2O.

Isotopic Abundance and Average Atomic Mass

• The average atomic mass is calculated using the masses and relative abundances of isotopes:

$\text{Average atomic mass} = \sum (\text{isotope mass} \times \text{fractional abundance})$

Example: If chlorine has two isotopes, Cl-35 (75.77%) and Cl-37 (24.23%), the average atomic mass is: $\text{Average atomic mass} = (0.7577 \times 34.969) + (0.2423 \times 36.966) = 35.45$ u$

Electron Configuration

• Describes the arrangement of electrons in an atom.

• Follows the Aufbau principle, Pauli exclusion principle, and Hund's rule.

Example: The electron configuration of oxygen (O, Z=8) is 1s2 2s2 2p4.

Unit 2: Molecular and Ionic Compound Structure and Properties

Lewis Structures and Molecular Geometry

Lewis structures show the arrangement of valence electrons in molecules. Molecular geometry describes the 3D arrangement of atoms.

• Electron domain geometry considers all electron pairs (bonding and lone pairs).

• Molecular geometry considers only the arrangement of atoms.

• Bond angle is the angle between adjacent bonds.

Example: CO2 has a linear geometry with a bond angle of 180°.

Polarity and Electronegativity

• Electronegativity is the ability of an atom to attract electrons in a bond.

• A bond is polar if there is a significant difference in electronegativity between the atoms.

Example: HCl is polar because Cl is more electronegative than H.

Unit 3: Intermolecular Forces and Properties

Types of Intermolecular Forces

• London Dispersion Forces (LDF): Weak, present in all molecules, increase with molecular size.

• Dipole-Dipole Forces: Occur between polar molecules.

• Hydrogen Bonding: Strong dipole-dipole interaction involving H bonded to N, O, or F.

Example: Water (H2O) exhibits hydrogen bonding.

Force	Relative Strength	Relative Boiling Point	Polarity	Example
London Dispersion	Weakest	Lowest	Nonpolar	CH4
Dipole-Dipole	Intermediate	Intermediate	Polar	HCl
Hydrogen Bonding	Strongest	Highest	Polar (with H-F, H-O, H-N)	H2O

Properties of Solids, Liquids, and Gases

• Solids: Definite shape and volume, particles in fixed positions.

• Liquids: Definite volume, indefinite shape, particles can move past each other.

• Gases: Indefinite shape and volume, particles move freely.

Unit 4: Chemical Reactions and Stoichiometry

Types of Chemical Reactions

• Synthesis: $A + B \rightarrow AB$

• Decomposition: $AB \rightarrow A + B$

• Single Replacement: $A + BC \rightarrow AC + B$

• Double Replacement: $AB + CD \rightarrow AD + CB$

• Combustion: Hydrocarbon + O2 $\rightarrow$ CO2 + H2O

Balancing Chemical Equations

• Ensure the same number of each atom on both sides of the equation.

Example: $2H_2 + O_2 \rightarrow 2H_2O$

Stoichiometry

• Calculations based on balanced chemical equations to determine the amounts of reactants and products.

• Use mole ratios from the coefficients in the balanced equation.

Example: How many grams of CO2 are produced from 10.0 g of propane (C3H8) combusted in excess oxygen?

Unit 5: Kinetics

Reaction Rates and Rate Laws

• Reaction rate is the change in concentration of a reactant or product per unit time.

• Rate law expresses the rate as a function of reactant concentrations:

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

• Order of reaction is the sum of the exponents in the rate law.

Factors Affecting Reaction Rates

• Concentration of reactants

• Temperature

• Catalysts

• Surface area

Activation Energy and Reaction Mechanisms

• Activation energy (Ea) is the minimum energy required for a reaction to occur.

• Catalysts lower the activation energy, increasing the reaction rate.

• Reaction mechanism is the sequence of elementary steps that make up a complex reaction.

Example: For a two-step mechanism, the slowest step determines the rate law.

Half-Life and First-Order Reactions

• The half-life ($t_{1/2}$) of a first-order reaction is the time required for half the reactant to be consumed:

$t_{1/2} = \frac{0.693}{k}$

Additional info:

• This review covers topics from Ch.1 (Introduction), Ch.2 (Atoms, Molecules, and Ions), Ch.3 (Chemical Reactions and Stoichiometry), Ch.4 (Reactions in Aqueous Solution), Ch.5 (Thermochemistry), Ch.6 (Electronic Structure of Atoms), Ch.7 (Periodic Properties), Ch.8 (Bonding), Ch.9 (Molecular Geometry), Ch.10 (Gases), Ch.11 (Liquids and Intermolecular Forces), Ch.12 (Solids), Ch.13 (Solutions), and Ch.14 (Kinetics).

• Tables and diagrams referenced in the questions are used for classification, comparison, and calculation practice.