General Chemistry Practice Exam Study Guide

General Chemistry Fundamentals

Chemical Tools: Experimentation & Measurement

Understanding measurement, significant figures, and units is essential in chemistry for accurate data collection and analysis.

• Significant Figures: The number of digits in a measurement that are known with certainty plus one estimated digit. For example, a ruler measurement should be reported with the correct number of significant figures.

• Density: Defined as mass per unit volume. The formula is: $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

• Accuracy vs. Precision: Accuracy refers to how close a measurement is to the true value, while precision refers to how close repeated measurements are to each other.

• Units and Conversions: Common units include grams (g), liters (L), and meters (m). Always use dimensional analysis for conversions.

• Example: Measuring the length of a metal bar using a ruler and reporting the value with the correct significant figures.

Atoms, Molecules & Ions

Chemistry is the study of matter, which is composed of atoms, molecules, and ions. Understanding their structure and properties is foundational.

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

• Isotopes: Atoms of the same element with different numbers of neutrons. The average atomic mass is calculated using isotopic abundances: $\text{Average Atomic Mass} = \sum (\text{Fractional Abundance} \times \text{Isotope Mass})$

• Ions: Atoms or molecules that have gained or lost electrons, resulting in a charge.

• Molecules: Two or more atoms bonded together.

• Example: Determining the number of protons, neutrons, and electrons in an isotope of potassium, $^{39}_{19}\text{K}$.

Mass Relationships in Chemical Reactions

Stoichiometry involves the calculation of reactants and products in chemical reactions using balanced equations.

• Mole Concept: The mole is a counting unit in chemistry. $1$mole$= 6.022 \times 10^{23}$ particles.

• Balancing Equations: Ensures the conservation of mass and atoms in a reaction.

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

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

• Example: Calculating the mass of $\text{CO}_2$ produced from the reaction of $\text{C}_2\text{H}_6$ with $\text{O}_2$.

Lab Techniques and Procedures

Proper laboratory techniques are crucial for safety and accuracy in chemical experimentation.

• Measurement: Use appropriate tools (balances, graduated cylinders) and report values with correct significant figures.

• Separation Techniques: Filtration, distillation, and chromatography are common methods to separate mixtures.

• Example: Determining the density of a metal by measuring its mass and volume.

Classification of Matter

Pure Substances and Mixtures

Matter can be classified as pure substances (elements and compounds) or mixtures (homogeneous and heterogeneous).

• Element: A substance composed of only one type of atom.

• Compound: A substance composed of two or more elements chemically bonded.

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

• Heterogeneous Mixture: Non-uniform composition (e.g., salad).

• Example: Air is a homogeneous mixture; granite is a heterogeneous mixture.

Physical and Chemical Properties

Properties of matter are classified as physical or chemical, and as intensive or extensive.

• Physical Property: Can be observed without changing the substance (e.g., melting point).

• Chemical Property: Describes the ability of a substance to undergo chemical change (e.g., flammability).

• Intensive Property: Independent of the amount of substance (e.g., density).

• Extensive Property: Depends on the amount of substance (e.g., mass).

• Example: Boiling water is a physical change; rusting iron is a chemical change.

Periodic Table & Atomic Structure

Periodic Table Organization

The periodic table arranges elements by increasing atomic number and groups elements with similar properties.

• Groups/Families: Vertical columns with similar chemical properties (e.g., Group 7A: Halogens).

• Periods: Horizontal rows.

• Metals, Nonmetals, Metalloids: Metals are typically shiny and conductive; nonmetals are varied; metalloids have intermediate properties.

• Example: The shaded area in a periodic table may indicate metals.

Electronic Structure of Atoms

Electrons are arranged in shells and subshells according to quantum numbers and the Aufbau principle.

• Quantum Numbers: Describe the energy, shape, and orientation of atomic orbitals.

• Electron Configuration: Shows the arrangement of electrons in an atom. Example for Iodine: $1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2 4d^{10} 5p^5$

• Hund's Rule: Electrons fill degenerate orbitals singly before pairing.

• Example: Identifying the ground state electron configuration for Ge.

Chemical Bonding

Ionic and Covalent Bonding

Chemical bonds form between atoms to achieve stability. Ionic bonds involve electron transfer; covalent bonds involve electron sharing.

• Ionic Bond: Formed between metals and nonmetals; involves transfer of electrons.

• Covalent Bond: Formed between nonmetals; involves sharing of electrons.

• Lewis Structures: Diagrams showing the arrangement of electrons in molecules.

• Resonance Structures: Multiple valid Lewis structures for a molecule.

• Example: Drawing the Lewis structure for $\text{OCl}_2$ and identifying resonance forms for $\text{NCO}^-$.

Bond Polarity and Molecular Geometry

The shape and polarity of molecules affect their physical and chemical properties.

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

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

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

• Example: Determining the geometry of $\text{SF}_4$ and the hybridization of carbon in $\text{CH}_2\text{O}$.

Thermochemistry

Energy Changes in Chemical Reactions

Chemical reactions involve energy changes, typically in the form of heat (enthalpy).

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

• Endothermic vs. Exothermic: Endothermic reactions absorb heat; exothermic reactions release heat.

• Calorimetry: Measurement of heat changes using a calorimeter.

• Work and Internal Energy: $\Delta E = q + w$ where $q$ is heat and $w$ is work.

• Example: Calculating $\Delta H$ for the combustion of methane and determining the heat capacity of a metal.

Gases: Properties and Behavior

Gas Laws

Gases follow specific 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: $PV = nRT$

• Standard Temperature and Pressure (STP): $0^\circ C$ and $1$ atm

• Example: Calculating the volume of $1$ mol of gas at STP.

Solutions and Their Properties

Types of Solutions and Concentration

Solutions are homogeneous mixtures of solute and solvent. Concentration is commonly expressed as molarity.

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

• Preparation of Solutions: Calculating the mass of solute needed for a given volume and concentration.

• Example: Determining the grams of $\text{KNO}_3$ needed to make $250$mL of $0.133$ M solution.

Tables

Threshold Frequencies for Metals (Photoelectric Effect)

The table below lists the threshold frequencies for different metals, which are important in understanding the photoelectric effect.

Metal	v, Hz
Cu	4.751 x 1014
Na	5.205 x 1014
Ag	6.626 x 1014
Sn	6.937 x 1014
Al	7.757 x 1014
Au	8.205 x 1014
Pb	8.625 x 1014
Zn	1.006 x 1015
Fe	1.122 x 1015
Pt	1.187 x 1015
Ca	1.320 x 1015
Mg	1.382 x 1015

Additional info:

• This study guide covers topics from Ch.1 (Measurement), Ch.2 (Atoms, Molecules & Ions), Ch.3 (Stoichiometry), Ch.4 (Reactions in Aqueous Solution), Ch.5 (Electronic Structure), Ch.6-8 (Bonding), Ch.9 (Thermochemistry), Ch.10 (Gases), Ch.13 (Solutions), and includes lab techniques and mathematical operations.

• Some questions also address periodic trends, molecular geometry, and intermolecular forces.