Comprehensive Study Guide for CHEM 106 Final Exam Topics

Basic Chemistry

Elements, Compounds, Mixtures, and Particulate Nature of Matter

Understanding the foundational concepts of chemistry is essential for further study. Matter is composed of atoms, which combine to form elements, compounds, and mixtures.

• Element: A pure substance consisting of only one type of atom.

• Compound: A substance formed from two or more elements chemically bonded in fixed proportions.

• Mixture: A combination of two or more substances not chemically bonded.

• Particulate Nature of Matter: Matter is made up of tiny particles (atoms, molecules, ions).

• Example: Water (H2O) is a compound; air is a mixture.

Atomic Structure

Isotopes, Atomic Number, Mass Number, Atomic Mass

Atoms consist of protons, neutrons, and electrons. Their arrangement determines the atom's identity and properties.

• 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.

• Atomic Mass: Weighted average mass of an element's isotopes.

• Example: Carbon-12 and Carbon-14 are isotopes of carbon.

Electrons in Atoms

Orbitals, Quantum Numbers, Electron Configuration, Electromagnetic Spectrum

Electrons occupy specific energy levels and orbitals around the nucleus, described by quantum numbers.

• Quantum Numbers: Describe the properties of atomic orbitals and electrons.

• Electron Configuration: Distribution of electrons among orbitals.

• Electromagnetic Spectrum: Range of all types of electromagnetic radiation.

• Example: The electron configuration of oxygen: 1s2 2s2 2p4.

Periodic Trends

Electron Affinity, Electronegativity, Ionization Energy, Ionic Radius

Periodic trends describe how certain properties of elements change across the periodic table.

• Ionization Energy: Energy required to remove an electron from an atom.

• Electronegativity: Ability of an atom to attract electrons in a bond.

• Electron Affinity: Energy change when an atom gains an electron.

• Ionic Radius: Size of an ion compared to its neutral atom.

• Example: Ionization energy increases across a period and decreases down a group.

Chemical Bonding

Lewis Dot Diagrams, VSEPR, Bond Energy, Bond Types

Chemical bonds form when atoms share or transfer electrons. The type of bond affects molecular properties.

• Lewis Dot Diagrams: Visual representation of valence electrons.

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

• Bond Types: Ionic, covalent, polar covalent, metallic.

• Bond Energy: Energy required to break a bond.

• Example: Water is a polar covalent molecule with bent geometry.

Chemical Quantities

Mole, Molar Mass, Empirical Formula, Stoichiometry

Quantitative chemistry involves measuring and calculating amounts of substances in reactions.

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

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

• Empirical Formula: Simplest whole-number ratio of elements in a compound.

• Stoichiometry: Calculation of reactants and products in chemical reactions.

• Example: $\text{C}_6\text{H}_{12}\text{O}_6$ has an empirical formula of $\text{CH}_2\text{O}$.

Types of Chemical Reactions

Classification, Balancing, Net Ionic Equations

Chemical reactions are classified by the changes that occur and must be balanced to obey the law of conservation of mass.

• Types: Synthesis, decomposition, single replacement, double replacement, combustion.

• Balancing: Ensures equal numbers of each atom on both sides of the equation.

• Net Ionic Equations: Show only the species that change during the reaction.

• Example: $2\text{H}_2 + \text{O}_2 \rightarrow 2\text{H}_2\text{O}$

Solution Chemistry

Types of Solutions, Solubility, Concentration, Solution Stoichiometry

Solutions are homogeneous mixtures of solute and solvent. Their properties depend on concentration and solubility.

• Solubility: Maximum amount of solute that can dissolve in a solvent at a given temperature.

• Concentration: Amount of solute per unit volume of solution (e.g., molarity).

• Solution Stoichiometry: Calculations involving reactions in solution.

• Example: Molarity ($M$) is calculated as $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

Gases

Gas Laws, Relationships, Properties

Gases have unique properties described by several laws relating pressure, volume, temperature, and amount.

• Boyle's Law: $P_1V_1 = P_2V_2$ (at constant temperature)

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

• Ideal Gas Law: $PV = nRT$

• Example: Doubling the pressure halves the volume (Boyle's Law).

Thermochemistry

Heat, Enthalpy, Calorimetry

Thermochemistry studies energy changes in chemical reactions, especially heat transfer.

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

• Calorimetry: Measurement of heat changes.

• Example: $q = mC\Delta T$ (heat = mass × specific heat × temperature change)

Thermodynamics

Entropy, Gibbs Free Energy

Thermodynamics explains the direction and spontaneity of chemical processes.

• Entropy ($S$): Measure of disorder in a system.

• Gibbs Free Energy ($\Delta G$): Determines spontaneity: $\Delta G = \Delta H - T\Delta S$

• Example: A negative $\Delta G$ indicates a spontaneous reaction.

Kinetics

Rates of Reaction, Rate Laws, Activation Energy

Chemical kinetics studies the speed of reactions and the factors that affect them.

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

• Activation Energy: Minimum energy required for a reaction to occur.

• Half-life ($t_{1/2}$): Time required for half the reactant to be consumed.

• Example: First-order reaction: $[A]_t = [A]_0 e^{-kt}$

Equilibrium

Equilibrium Constant, Le Châtelier's Principle

Chemical equilibrium occurs when the rates of forward and reverse reactions are equal.

• Equilibrium Constant ($K$): Ratio of product concentrations to reactant concentrations at equilibrium.

• Le Châtelier's Principle: System at equilibrium responds to disturbances to restore balance.

• Example: $K = \frac{[C]^c[D]^d}{[A]^a[B]^b}$ for $aA + bB \rightleftharpoons cC + dD$

Acids and Bases

Definitions, pH, Titrations, Buffers

Acids and bases are defined by their ability to donate or accept protons. Their strength and concentration are measured by pH.

• Arrhenius Definition: Acids produce H+, bases produce OH- in water.

• Bronsted-Lowry: Acids donate protons, bases accept protons.

• pH: $pH = -\log[H^+]$

• Titration: Technique to determine concentration of an acid or base.

• Buffer: Solution that resists changes in pH.

• Example: Vinegar (acetic acid) is a weak acid.

Redox Reactions

Oxidation, Reduction, Balancing, Electrochemistry

Redox reactions involve the transfer of electrons between species, affecting oxidation states.

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Electrochemistry: Study of chemical processes that cause electrons to move.

• Example: $\text{Zn} + \text{Cu}^{2+} \rightarrow \text{Zn}^{2+} + \text{Cu}$

Radioactivity and Nuclear Chemistry

Types of Decay, Half-life, Nuclear Reactions

Nuclear chemistry focuses on changes in atomic nuclei, including radioactive decay.

• Alpha Decay: Emission of an alpha particle ($^4_2\text{He}$).

• Beta Decay: Emission of a beta particle (electron or positron).

• Gamma Decay: Emission of gamma radiation (high-energy photons).

• Half-life: Time for half of a radioactive sample to decay.

• Example: $^{14}\text{C} \rightarrow ^{14}\text{N} + e^-$ (beta decay)

Colligative Properties

Properties of Solutions

Colligative properties depend on the number of solute particles in a solution, not their identity.

• Boiling Point Elevation: $\Delta T_b = iK_b m$

• Freezing Point Depression: $\Delta T_f = iK_f m$

• Osmotic Pressure: $\Pi = iMRT$

• Example: Salt added to water lowers its freezing point.