General Chemistry I: Comprehensive Study Guide (Practice Final Exam Topics)

General Chemistry I: Comprehensive Study Guide

Introduction

This study guide covers foundational topics in General Chemistry, as reflected in a comprehensive practice final exam. The material spans atomic structure, chemical reactions, stoichiometry, gases, thermochemistry, quantum mechanics, periodic trends, bonding, molecular structure, solutions, kinetics, equilibrium, acids and bases, and more. Each section below summarizes key concepts, definitions, formulas, and example applications relevant to first-semester college chemistry.

Atoms, Elements, and Compounds

Classification of Matter

• Element: A pure substance consisting of only one type of atom (e.g., Carbon).

• Compound: A substance composed of two or more elements chemically combined in fixed proportions (e.g., Water (H2O)).

• Mixture: A physical blend of two or more substances that are not chemically combined (e.g., Air).

Example: Classify the following: Carbon Dioxide (compound), Air (mixture), Water (compound).

Atomic Structure

• Protons: Positively charged particles in the nucleus.

• Neutrons: Neutral particles in the nucleus.

• Electrons: Negatively charged particles in orbitals around the nucleus.

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

Example: 35Cl and 37Cl are isotopes of chlorine.

Electron Configuration

• Describes the arrangement of electrons in an atom's orbitals.

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

Example: The electron configuration of sodium (Na): $1s^2 2s^2 2p^6 3s^1$

Chemical Reactions and Stoichiometry

Types of Chemical Reactions

• Synthesis: Two or more substances combine to form one product.

• Decomposition: A compound breaks down into two or more simpler substances.

• Single Replacement: One element replaces another in a compound.

• Double Replacement: Exchange of ions between two compounds.

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

Balancing Chemical Equations

• Law of Conservation of Mass: Matter is neither created nor destroyed.

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

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

Stoichiometry

• Relates quantities of reactants and products using balanced equations.

• Mole: The SI unit for amount of substance; $1$mole = $6.022 \times 10^{23}$ particles (Avogadro's number).

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

Example: Calculate the mass of $1.00 \times 10^{23}$ molecules of water: $\text{Moles} = \frac{1.00 \times 10^{23}}{6.022 \times 10^{23}}$ $\text{Mass} = \text{Moles} \times \text{Molar Mass}$

Chemical Quantities and Unit Conversions

Unit Conversions

• Use conversion factors to change units (e.g., cm to km, mi to m).

• Dimensional analysis ensures correct unit cancellation.

Example: Convert $2.18 \times 10^5$ cm to km: $2.18 \times 10^5 \text{ cm} \times \frac{1 \text{ m}}{100 \text{ cm}} \times \frac{1 \text{ km}}{1000 \text{ m}} = 2.18 \text{ km}$

Significant Figures

• Indicate the precision of a measured or calculated quantity.

• Rules for counting significant figures depend on zeros and decimal placement.

Example: The number 0.003450 has 4 significant figures.

Gases and Gas Laws

Properties of Gases

• Gases have variable shape and volume, low density, and are compressible.

Gas 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)

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

• Ideal Gas Law: $PV = nRT$

Example: Calculate the volume of a gas at STP given moles using $V = nRT/P$.

Thermochemistry

Energy and Heat

• Energy: The capacity to do work or transfer heat.

• Heat (q): Energy transferred due to temperature difference.

• Work (w): Energy used to move an object against a force.

First Law of Thermodynamics

• Energy is conserved: $\Delta E = q + w$

Quantum Mechanics and Atomic Structure

Quantum Numbers

• Principal (n): Energy level (n = 1, 2, 3, ...)

• Angular Momentum (l): Sublevel (l = 0 to n-1)

• Magnetic (ml): Orientation (-l to +l)

• Spin (ms): Electron spin (+1/2 or -1/2)

Example: For a 3p orbital: n = 3, l = 1, ml = -1, 0, or +1, ms = +1/2 or -1/2

Electromagnetic Radiation

• Energy of a photon: $E = h\nu$

• Relationship between wavelength and frequency: $c = \lambda \nu$

Example: Calculate the wavelength of light with frequency $3.22 \times 10^{15} \text{ s}^{-1}$: $\lambda = \frac{c}{\nu}$

Periodic Properties of the Elements

Trends in the Periodic Table

• 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.

• Electron Affinity: Energy change when an electron is added to an atom.

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

Example: Arrange Cl, S, Se, Ge in order of increasing ionization energy: Ge < Se < S < Cl

Chemical Bonding and Molecular Structure

Ionic and Covalent Bonds

• Ionic Bond: Transfer of electrons from metal to nonmetal, forming ions.

• Covalent Bond: Sharing of electrons between nonmetals.

Example: NaCl is ionic; H2O is covalent.

Lewis Structures

• Show arrangement of valence electrons in molecules.

• Octet rule: Atoms tend to have 8 electrons in their valence shell.

Example: Draw Lewis structures for SO3 with and without expanded octet.

VSEPR Theory and Molecular Geometry

• Predicts 3D shape of molecules based on electron domain repulsion.

• Common geometries: linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.

Example: CH4 is tetrahedral; CO2 is linear.

States of Matter and Intermolecular Forces

Liquids, Solids, and Gases

• Solid: Definite shape and volume.

• Liquid: Definite volume, variable shape.

• Gas: Variable shape and volume.

Intermolecular Forces (IMF)

• London Dispersion: Weakest, present in all molecules.

• Dipole-Dipole: Between polar molecules.

• Hydrogen Bonding: Strongest, occurs when H is bonded to N, O, or F.

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

Solutions and Concentrations

Types of Solutions

• Solute: Substance dissolved.

• Solvent: Substance doing the dissolving.

Concentration Units

• Molarity (M): $M = \frac{\text{moles solute}}{\text{liters solution}}$

• Percent by mass, volume, or mole fraction

Example: Calculate the molarity of a solution made by dissolving 17.4 g Na2CrO4 in 500 mL of solution.

Acids, Bases, and Chemical Equilibrium

Acids and Bases

• Arrhenius Acid: Produces H+ in water.

• Arrhenius Base: Produces OH- in water.

• Bronsted-Lowry Acid/Base: Proton donor/acceptor.

Chemical Equilibrium

• At equilibrium, the rate of forward and reverse reactions are equal.

• Equilibrium constant ($K$) expresses the ratio of product to reactant concentrations at equilibrium.

Example: $K = \frac{[\text{products}]}{[\text{reactants}]}$

Sample Table: Comparison of Bond Types

Bond Type	Formation	Example	Properties
Ionic	Transfer of electrons	NaCl	High melting point, conducts electricity when molten
Covalent	Sharing of electrons	H2O	Low melting point, does not conduct electricity
Metallic	Delocalized electrons	Fe	Conducts electricity, malleable

Additional Topics

• Empirical and Molecular Formulas: Empirical is the simplest ratio; molecular is the actual number of atoms.

• Limiting Reactant: The reactant that is completely consumed first in a reaction.

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

• Colligative Properties: Properties that depend on the number of solute particles (e.g., boiling point elevation, freezing point depression).

Conclusion

This guide summarizes the essential concepts and problem types encountered in a first-semester General Chemistry course. Mastery of these topics will prepare students for comprehensive exams and further study in chemistry.