General Chemistry Final Exam Study Guide: Key Concepts and Practice

Atoms & Elements

Atomic Structure and Properties

Atoms are the fundamental units of matter, composed of protons, neutrons, and electrons. Their arrangement and number determine the element's identity and chemical behavior.

• Atomic Number (Z): Number of protons in the nucleus; defines the element.

• Mass Number (A): Sum of protons and neutrons.

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

• Electron Configuration: Distribution of electrons among atomic orbitals, e.g., $1s^2 2s^2 2p^6$.

• Valence Electrons: Electrons in the outermost shell, crucial for chemical bonding.

• Periodic Table: Organizes elements by increasing atomic number and recurring chemical properties.

Example: Sodium (Na) has atomic number 11, electron configuration $1s^2 2s^2 2p^6 3s^1$.

Chemical Reactions

Types and Balancing of Chemical Reactions

Chemical reactions involve the rearrangement of atoms to form new substances. They are represented by balanced chemical equations.

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

• Balancing Equations: Ensures conservation of mass and charge. Coefficients are adjusted so that the number of atoms of each element is equal on both sides.

• Oxidation and Reduction: Oxidation is loss of electrons; reduction is gain of electrons. The oxidizing agent is reduced, and the reducing agent is oxidized.

Example: Photosynthesis: $6CO_2 + 6H_2O ightarrow C_6H_{12}O_6 + 6O_2$

Chemical Quantities & Aqueous Reactions

Stoichiometry and Solution Chemistry

Stoichiometry involves quantitative relationships in chemical reactions, including mole calculations and limiting reactants.

• Mole Concept: 1 mole = $6.022 \times 10^{23}$ particles (Avogadro's number).

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

• Limiting Reactant: The reactant that is completely consumed first, limiting the amount of product formed.

• Concentration: Amount of solute per unit volume of solution, commonly expressed as molarity ($M = \frac{\text{mol solute}}{\text{L solution}}$).

• Solution Types: Homogeneous (uniform composition) and heterogeneous (non-uniform composition).

Example: Calculating moles of NaCl in 0.0886 g: $\text{mol NaCl} = \frac{0.0886\,g}{58.44\,g/mol}$

Gases

Gas Laws and Properties

Gases are described by several laws relating pressure, volume, temperature, and amount.

• Ideal Gas Law: $PV = nRT$

• STP Conditions: Standard Temperature and Pressure (0°C, 1 atm).

• Partial Pressure: Pressure exerted by each gas in a mixture.

• Density of Gases: $\text{Density} = \frac{PM}{RT}$

Example: Volume of 4.5 L gas at STP: $n = \frac{PV}{RT}$

Thermochemistry

Energy Changes in Chemical Reactions

Thermochemistry studies the heat involved in chemical processes.

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

• Specific Heat Capacity ($c$): Amount of heat required to raise the temperature of 1 g of substance by 1°C.

• Calorimetry: Measurement of heat changes.

• Endothermic vs. Exothermic: Endothermic absorbs heat; exothermic releases heat.

Example: $q = mc\Delta T$

Quantum Mechanics

Atomic Orbitals and Electron Transitions

Quantum mechanics explains the behavior of electrons in atoms.

• Principal Quantum Number ($n$): Indicates energy level.

• Electron Transitions: Electrons absorb or emit energy to move between levels.

• Electromagnetic Spectrum: Light is emitted or absorbed during transitions.

Example: $E = h\nu$ where $h$ is Planck's constant and $\nu$ is frequency.

Periodic Properties of the Elements

Trends in the Periodic Table

Elements show periodic trends in properties such as electronegativity, ionization energy, and atomic radius.

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

• Ionization Energy: Energy required to remove an electron; increases across a period.

• Atomic Radius: Size of an atom; decreases across a period, increases down a group.

Example: Fluorine is the most electronegative element.

Bonding & Molecular Structure

Covalent, Ionic, and Metallic Bonds

Chemical bonds form between atoms to create molecules and compounds.

• Ionic Bonds: Transfer of electrons between metals and nonmetals.

• Covalent Bonds: Sharing of electrons between nonmetals.

• Metallic Bonds: Delocalized electrons among metal atoms.

• Bond Polarity: Difference in electronegativity leads to polar or nonpolar bonds.

Example: NaCl is ionic; H2O is polar covalent.

Molecular Shapes & Valence Bond Theory

VSEPR and Hybridization

The shape of molecules is predicted by the Valence Shell Electron Pair Repulsion (VSEPR) theory and hybridization of atomic orbitals.

• VSEPR: Electron pairs around a central atom arrange to minimize repulsion.

• Common Geometries: Linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.

• Hybridization: Mixing of atomic orbitals to form new hybrid orbitals (e.g., $sp^3$, $sp^2$, $sp$).

Example: Methane (CH4) is tetrahedral ($sp^3$ hybridization).

Liquids, Solids & Intermolecular Forces

States of Matter and Forces Between Molecules

Matter exists as solids, liquids, and gases, with intermolecular forces determining physical properties.

• Types of Forces: London dispersion, dipole-dipole, hydrogen bonding.

• Properties: Boiling point, melting point, solubility depend on intermolecular forces.

Example: Water has strong hydrogen bonding, leading to high boiling point.

Solutions

Solution Composition and Concentration

Solutions are homogeneous mixtures of solute and solvent.

• Molarity ($M$): $M = \frac{\text{mol solute}}{\text{L solution}}$

• Mass Percent: $\text{Mass percent} = \frac{\text{mass of solute}}{\text{mass of solution}} \times 100\%$

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

Example: Calculating molarity of NaCl in water.

Chemical Kinetics

Rates of Chemical Reactions

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

• Rate: Change in concentration of reactants or products per unit time.

• Factors: Concentration, temperature, catalysts, surface area.

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

Chemical Equilibrium

Dynamic Balance in Chemical Systems

At equilibrium, the rates of forward and reverse reactions 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.

Example: $N_2 + 3H_2 \rightleftharpoons 2NH_3$

Acid and Base Equilibrium

Properties and Calculations of Acids and Bases

Acids donate protons (H+), bases accept protons. Their strength is measured by dissociation in water.

• Strong vs. Weak Acids/Bases: Strong acids/bases dissociate completely; weak only partially.

• pH Calculation: $pH = -\log[H^+]$

• Buffer Solutions: Resist changes in pH upon addition of acid or base.

Example: Hydrochloric acid (HCl) is a strong acid; acetic acid (CH3COOH) is weak.

Chemical Thermodynamics

Energy, Entropy, and Free Energy

Thermodynamics studies energy changes and spontaneity of reactions.

• First Law: Energy is conserved.

• Second Law: Entropy of the universe increases in spontaneous processes.

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

Example: A reaction is spontaneous if $\Delta G < 0$.

Electrochemistry

Redox Reactions and Electrochemical Cells

Electrochemistry involves electron transfer and generation of electrical energy.

• Galvanic Cells: Convert chemical energy to electrical energy.

• Standard Electrode Potential ($E^\circ$): Measures tendency to gain electrons.

• Nernst Equation: $E = E^\circ - \frac{0.0592}{n} \log Q$

Example: Zn/Cu cell generates electricity via redox reaction.

Lab Techniques and Procedures

Measurement, Significant Figures, and Experimental Methods

Accurate measurement and proper lab techniques are essential in chemistry.

• Significant Figures: Digits that reflect the precision of a measurement.

• Dimensional Analysis: Method for converting units.

• Identifying Spectator Ions: Ions that do not participate in the chemical reaction.

Example: Calculating the number of significant figures in 0.010167 (5 significant figures).

Mathematical Operations and Functions

Calculations in Chemistry

Mathematical skills are required for quantitative chemical analysis.

• Unit Conversions: Converting between grams, moles, liters, etc.

• Using Formulas: Applying equations such as $PV = nRT$, $q = mc\Delta T$, $M = \frac{m}{n}$.

Example: Calculating mass from density and volume: $m = \text{density} \times \text{volume}$

Additional info:

This study guide covers a broad range of topics from a General Chemistry final exam, including atomic structure, chemical reactions, stoichiometry, gas laws, thermochemistry, quantum mechanics, periodic trends, bonding, molecular geometry, solutions, kinetics, equilibrium, acids and bases, thermodynamics, electrochemistry, lab techniques, and mathematical operations. Each topic is presented with definitions, key points, formulas, and examples to facilitate exam preparation.