General Chemistry Final Exam Multiple Choice Practice: Comprehensive Study Notes

Atoms & Elements

Atomic Structure and Electron Configuration

Atoms are the basic units of matter, composed of protons, neutrons, and electrons. The arrangement of electrons in shells and subshells determines chemical properties.

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

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

• Example: Sodium (Na): $1s^2 2s^2 2p^6 3s^1$; 1 valence electron.

Periodic Properties

Elements are arranged in the periodic table by increasing atomic number, revealing periodic trends.

• Electronegativity: Tendency to attract electrons; 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 has the highest electronegativity.

Chemical Reactions

Types of Chemical Reactions

Chemical reactions involve the rearrangement of atoms to form new substances.

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

• Decomposition: A compound breaks down into simpler substances.

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

• Double Displacement: Exchange of ions between two compounds.

• Example: $2KClO_3 \rightarrow 2KCl + 3O_2$ (decomposition)

Balancing Chemical Equations

Law of conservation of mass requires that the number of atoms of each element is the same on both sides of the equation.

• Steps: List elements, count atoms, adjust coefficients.

• Example: $4KNO_3 + 2K_2SO_4 + 2C \rightarrow 3K_2CO_3 + 3SO_2 + 2N_2$

Chemical Quantities & Aqueous Reactions

Mole Concept and Avogadro's Number

The mole is a fundamental unit for counting particles in chemistry.

• Avogadro's Number: $6.022 \times 10^{23}$ particles/mol.

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

• Example: 1 mole of NaCl = 58.44 g.

Stoichiometry

Stoichiometry involves calculations based on balanced chemical equations.

• Steps: Convert mass to moles, use mole ratios, convert back to mass or volume.

• Example: How many moles of $K_2SO_4$ are produced from 34.4 g BaCl$_2$?

Solutions and Concentration

Solutions are homogeneous mixtures; concentration is often expressed as molarity.

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

• Example: 0.5 mol NaCl in 1 L solution = 0.5 M NaCl.

Gases

Gas Laws

Gas behavior is described by several laws relating pressure, volume, temperature, and amount.

• Ideal Gas Law: $PV = nRT$

• Boyle's Law: $P_1V_1 = P_2V_2$ (at constant T)

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

• Example: Calculate volume of 1 mol gas at STP: $V = \frac{nRT}{P}$

Thermochemistry

Energy Changes in Reactions

Thermochemistry studies heat changes during chemical reactions.

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

• Specific Heat Capacity ($c$): $q = mc\Delta T$

• Example: Calculate heat required to raise temperature of 17.3 g substance by 57.0°C.

Quantum Mechanics

Electronic Structure and Light

Quantum mechanics explains atomic structure and the behavior of electrons.

• Energy Levels: Electrons occupy quantized energy levels.

• Photon Energy: $E = h\nu$

• Wavelength-Frequency Relationship: $c = \lambda \nu$

• Example: Calculate wavelength for frequency $4.72 \times 10^{14} \text{ s}^{-1}$.

Bonding & Molecular Structure

Covalent and Ionic Bonds

Bonds form when atoms share or transfer electrons to achieve stability.

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

• Covalent Bonds: Sharing of electrons between nonmetals.

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

• Example: NaCl is ionic; H$_2$O is polar covalent.

Molecular Geometry and VSEPR Theory

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

• Linear: 180° bond angle, e.g., CO$_2$

• Tetrahedral: 109.5° bond angle, e.g., CH$_4$

• Trigonal Planar: 120° bond angle, e.g., BF$_3$

• Example: SF$_6$ has octahedral geometry.

Solutions

Types of Mixtures

Mixtures can be homogeneous (solutions) or heterogeneous.

• Homogeneous Mixture: Uniform composition, e.g., salt water.

• Heterogeneous Mixture: Non-uniform, e.g., orange juice with pulp.

Solubility and Precipitation

Solubility is the ability of a substance to dissolve in a solvent. Precipitation occurs when solubility limits are exceeded.

• Example: Mixing AgNO$_3$ and NaCl forms a precipitate of AgCl.

Chemical Kinetics & Equilibrium

Reaction Rates

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

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

• Factors: Concentration, temperature, catalysts.

Chemical Equilibrium

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.

Acids, Bases, and Aqueous Equilibrium

Acid-Base Definitions

Acids donate protons (H$^+$), bases accept protons.

• Strong Acids/Bases: Completely ionize in water.

• Weak Acids/Bases: Partially ionize.

• pH Calculation: $\text{pH} = -\log[H^+]$

• Example: HCl is a strong acid; CH$_3$COOH is a weak acid.

Thermodynamics

First Law of Thermodynamics

Energy cannot be created or destroyed, only transformed.

• Internal Energy ($\Delta U$): $\Delta U = q + w$

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

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

• Example: Spontaneity determined by sign of $\Delta G$.

Electrochemistry

Redox Reactions

Electrochemistry studies electron transfer in chemical reactions.

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Oxidizing Agent: Causes oxidation, is reduced.

• Reducing Agent: Causes reduction, is oxidized.

• Example: $2KMnO_4 + 3Na_2SO_3 + H_2O \rightarrow 2MnO_2 + 3Na_2SO_4 + 2KOH$

Lab Techniques and Procedures

Measurement and Significant Figures

Accurate measurement and reporting are essential in chemistry.

• Significant Figures: Digits that reflect precision; rules for counting and rounding.

• Example: 0.010167 has 5 significant figures.

Density Calculations

Density is mass per unit volume.

• Formula: $\text{Density} = \frac{\text{mass}}{\text{volume}}$

• Example: Find density of a gold bar weighing 8.0 kg with volume 0.41 L.

Mathematical Operations and Functions

Unit Conversions and Calculations

Chemistry often requires converting between units and using dimensional analysis.

• Example: Convert grams to moles using molar mass.

• Example: Calculate number of molecules in a given mass.

HTML Table Example: Periodic Trends

Element	Electronegativity	Ionization Energy	Atomic Radius
F	Highest	High	Small
Na	Low	Low	Large
Cl	High	High	Medium

Additional info:

• Questions cover all major topics in General Chemistry, including atoms, bonding, reactions, stoichiometry, gases, thermochemistry, quantum mechanics, periodic properties, solutions, kinetics, equilibrium, acids/bases, thermodynamics, electrochemistry, and lab techniques.

• Some questions require calculation and application of formulas; others test conceptual understanding.

• Bonus section includes short answer and lab-based questions.