General Chemistry Final Exam Study Guide

Study of Matter

Phases and Classification of Matter

Matter exists in different phases and can be classified based on its physical and chemical properties.

• Phases of Matter: Solid, liquid, gas, and plasma.

• Classification: Elements, compounds, and mixtures (homogeneous and heterogeneous).

• Physical vs. Chemical Properties: Physical properties can be observed without changing the substance's identity (e.g., melting point, density), while chemical properties describe a substance's ability to undergo chemical changes (e.g., flammability).

Atoms, Elements, and Compounds

• Atoms: The basic unit of matter, consisting of protons, neutrons, and electrons.

• Elements: Pure substances consisting of only one type of atom.

• Compounds: Substances formed from two or more elements chemically combined in fixed ratios.

• Molecules: Two or more atoms bonded together.

Measurement and Units

• SI Units: Standard units for scientific measurements (meter, kilogram, second, mole, etc.).

• Temperature Scales: Celsius (°C), Fahrenheit (°F), Kelvin (K).

• Density: $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

Atomic Theory

Atomic Structure and Electron Configuration

Atomic theory explains the structure of atoms and the arrangement of electrons.

• 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.

• Electron Configuration: Distribution of electrons in atomic orbitals.

Periodic Table

• Groups: Vertical columns with similar chemical properties.

• Periods: Horizontal rows indicating energy levels.

• Periodic Trends: Atomic radius, ionization energy, electron affinity, electronegativity.

Chemical Bonding and Structure

Types of Chemical Bonds

• Ionic Bonds: Transfer of electrons from metal to nonmetal.

• Covalent Bonds: Sharing of electrons between nonmetals.

• Metallic Bonds: Delocalized electrons among metal atoms.

Lewis Structures and Molecular Geometry

• Lewis Structures: Diagrams showing valence electrons and bonding.

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

Chemical Reactions

Types of Chemical Reactions

Chemical reactions can be classified based on the changes that occur.

• Combination (Synthesis) Reaction: Two or more substances combine to form one product.

• Decomposition Reaction: A single compound breaks down into two or more products.

• Single-Replacement Reaction: An element replaces another in a compound.

• Double-Replacement Reaction: Exchange of ions between two compounds.

• Combustion Reaction: A substance reacts with oxygen, releasing energy.

Balancing Chemical Equations

• Law of Conservation of Mass: The total mass of reactants equals the total mass of products.

• Coefficients are used to balance the number of atoms of each element on both sides of the equation.

Stoichiometry and Molar Ratios

Mole Concept

• Mole: The amount of substance containing $6.022 \times 10^{23}$ entities (Avogadro's number).

• Molar Mass: Mass of one mole of a substance (g/mol).

Stoichiometric Calculations

• Use balanced chemical equations to relate moles of reactants and products.

• Theoretical Yield: Maximum amount of product that can be formed.

• Percent Yield: $\text{Percent Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100\%$

Gases

Gas Laws

Gas behavior is described by several fundamental laws.

• 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)

• Gay-Lussac's Law: $\frac{P_1}{T_1} = \frac{P_2}{T_2}$ (at constant V and n)

• Avogadro's Law: $\frac{V_1}{n_1} = \frac{V_2}{n_2}$ (at constant P and T)

• Ideal Gas Law: $PV = nRT$

Gas Properties

• Relationship between pressure, volume, temperature, and amount of gas.

• Density and molar mass of gases.

Thermochemistry

Energy Changes in Chemical Reactions

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

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

• Enthalpy ($\Delta H$): Heat content of a system at constant pressure.

• Energy diagrams illustrate the energy changes during reactions.

Liquids, Solids, and Intermolecular Forces

Types of Intermolecular Forces

• London Dispersion Forces: Weak, temporary attractions in all molecules.

• Dipole-Dipole Forces: Attractions between polar molecules.

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

Properties of Liquids and Solids

• Melting and boiling points depend on intermolecular forces.

• Solubility, vapor pressure, and viscosity are influenced by molecular structure.

Acids and Bases

Definitions and Properties

• Arrhenius Acid: Produces H+ ions in water.

• Arrhenius Base: Produces OH- ions in water.

• Bronsted-Lowry Acid: Proton donor.

• Bronsted-Lowry Base: Proton acceptor.

• Lewis Acid: Electron pair acceptor.

• Lewis Base: Electron pair donor.

Acid-Base Reactions and Calculations

• Neutralization: Acid reacts with base to form water and salt.

• Titration: Technique to determine concentration using a standard solution.

• Calculations using $C_1V_1 = C_2V_2$ for dilution and titration problems.

Redox Reactions

Oxidation and Reduction

• Oxidation: Loss of electrons (increase in oxidation number).

• Reduction: Gain of electrons (decrease in oxidation number).

• Assigning oxidation numbers to elements in compounds.

Sample Table: Types of Chemical Reactions

Type of Reaction	General Form	Example
Combination (Synthesis)	$A + B \rightarrow AB$	$2H_2 + O_2 \rightarrow 2H_2O$
Decomposition	$AB \rightarrow A + B$	$2H_2O \rightarrow 2H_2 + O_2$
Single-Replacement	$A + BC \rightarrow AC + B$	$Zn + 2HCl \rightarrow ZnCl_2 + H_2$
Double-Replacement	$AB + CD \rightarrow AD + CB$	$AgNO_3 + NaCl \rightarrow AgCl + NaNO_3$
Combustion	Hydrocarbon + $O_2 \rightarrow CO_2 + H_2O$	$CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O$