General Chemistry Study Guide: Core Concepts and Skills (Units 1–3)

Unit 1: Foundations of Chemistry

Properties of Matter

Understanding the properties of matter is fundamental in chemistry. Properties are classified as either intensive or extensive, and as physical or chemical.

• Intensive Properties: Do not depend on the amount of substance (e.g., density, boiling point).

• Extensive Properties: Depend on the amount of substance (e.g., mass, volume).

• Physical Properties: Can be observed without changing the substance's identity (e.g., color, melting point).

• Chemical Properties: Describe a substance's ability to undergo chemical changes (e.g., flammability, reactivity).

• Physical vs. Chemical Changes: Physical changes do not alter the chemical composition, while chemical changes result in new substances.

Example: Melting ice is a physical change; burning wood is a chemical change.

Types of Matter

Matter can be classified based on its composition and uniformity.

• Pure Substances: Have a fixed composition (elements and compounds).

• Mixtures: Combinations of two or more substances. Can be homogeneous (uniform, e.g., saltwater) or heterogeneous (non-uniform, e.g., salad).

• Separation Techniques: Methods such as filtration, distillation, and chromatography are used to separate mixtures based on physical properties.

Example: Salt can be separated from water by evaporation.

Ionic vs. Covalent Compounds

Chemical compounds are classified by the type of bonding between atoms.

• Ionic Compounds: Formed from metals and nonmetals; involve transfer of electrons.

• Covalent Compounds: Formed from nonmetals; involve sharing of electrons.

• Be able to write formulas using basic rules (e.g., binary ionic compounds).

Example: NaCl is ionic; H2O is covalent.

Significant Figures

Significant figures reflect the precision of a measurement.

• Use the correct number of significant figures in calculations.

• Understand rules for addition/subtraction and multiplication/division with significant figures.

• Know how to estimate uncertainty and distinguish between accuracy and precision.

Example: 2.50 × 3.0 = 7.5 (2 significant figures).

States of Matter

Matter exists in different physical states, each with unique properties.

• Solid: Definite shape and volume.

• Liquid: Definite volume, takes shape of container.

• Gas: No definite shape or volume; fills container.

Example: Water as ice (solid), liquid water, and steam (gas).

Density

Density is a measure of mass per unit volume.

• Formula: $d = \frac{m}{V}$

• Used to identify substances and solve problems involving mass and volume.

Example: If a block has a mass of 10 g and a volume of 2 cm3, its density is 5 g/cm3.

The Periodic Table

The periodic table organizes elements by increasing atomic number and properties.

• Metals: Conductive, malleable, shiny.

• Nonmetals: Poor conductors, brittle, dull.

• Metalloids: Properties intermediate between metals and nonmetals.

Dimensional Analysis

Dimensional analysis is a method for converting units using conversion factors.

• Set up conversion factors so units cancel appropriately.

• Useful for solving complex problems involving multiple unit conversions.

Example: To convert 10 inches to centimeters: $10 \text{ in} \times \frac{2.54 \text{ cm}}{1 \text{ in}} = 25.4 \text{ cm}$

Unit 2: Chemical Formulas and Solutions

Chemical Formulas

Chemical formulas represent the composition of compounds.

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

• Molecular Formula: Actual number of atoms of each element in a molecule.

• Hydrate Formula: Indicates water molecules associated with a compound (e.g., CuSO4·5H2O).

• Calculate percent composition and determine the amount of specific elements in a compound.

Example: The empirical formula of C6H12O6 is CH2O.

Properties of Solutions

Solutions are homogeneous mixtures of solute and solvent.

• Solute: Substance dissolved.

• Solvent: Substance doing the dissolving (often water).

• Understand solubility, miscibility, and the effect of temperature on solubility.

• Distinguish between saturated, unsaturated, and supersaturated solutions.

• Know which compounds are electrolytes (conduct electricity) and nonelectrolytes.

Example: Saltwater is a solution; NaCl is the solute, water is the solvent.

Colorimetry / Beer's Law

Colorimetry is used to determine the concentration of colored solutions using light absorption.

• Beer's Law: $A = \varepsilon l c$

• A: Absorbance, ε: Molar absorptivity, l: Path length, c: Concentration.

• Choose the best wavelength for maximum absorbance.

Molar Calculations

Mole calculations relate particles, mass, and chemical formulas.

• Use Avogadro's number ($6.022 \times 10^{23}$) to convert between moles and particles.

• Calculate molar mass from chemical formulas.

• Convert between mass, moles, and number of particles.

Example: 1 mole of H2O contains $6.022 \times 10^{23}$ molecules.

Acids

Acids are substances that increase the concentration of H+ ions in solution.

• pH is a measure of hydrogen ion concentration: $\text{pH} = -\log[H^+ ]$

• Be able to convert between pH and [H+].

Example: If [H+] = $1 \times 10^{-3}$ M, then pH = 3.

Unit 3: Atomic Structure and Nuclear Chemistry

Light Equations

Light exhibits both wave and particle properties, described by equations relating wavelength, frequency, and energy.

• Wavelength ($\lambda$), frequency ($\nu$), and energy ($E$) are related by:

  • $c = \lambda \nu$ (speed of light equation)

  • $E = h \nu$ (Planck's equation)

• Electrons absorb or emit energy as light when they move between energy levels.

Example: Calculate the energy of a photon with a frequency of $5.0 \times 10^{14}$ Hz: $E = (6.626 \times 10^{-34} \text{ J·s})(5.0 \times 10^{14} \text{ Hz})$

Nuclear Chemistry and Decay

Nuclear chemistry studies changes in atomic nuclei, including radioactive decay.

• Types of decay: alpha ($\alpha$), beta ($\beta$), electron capture, positron emission.

• Balance nuclear equations for decay processes.

• Calculate half-life and remaining quantity of radioactive isotopes.

• Understand fusion (combining nuclei) and fission (splitting nuclei).

Example: $^{14}_{6}\text{C} \rightarrow ^{14}_{7}\text{N} + \beta^-$ (beta decay)

Atomic Structure

Atoms consist of protons, neutrons, and electrons.

• Atomic number = number of protons.

• Mass number = protons + neutrons.

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

• Calculate average atomic mass using isotopic abundances.

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

Atomic History

The atomic model has evolved through experiments and discoveries.

• Key contributors: Dalton (atomic theory), Thomson (electron), Millikan (electron charge), Rutherford (nucleus), Bohr (energy levels).

• Modern atomic theory incorporates quantum mechanics.

Electron Configuration

Electron configuration describes the arrangement of electrons in an atom.

• Electrons fill orbitals in order of increasing energy (Aufbau principle).

• Use the periodic table to determine configurations for elements and ions.

• Recognize orbital diagrams and predict chemical behavior based on valence electrons.

Example: The electron configuration of oxygen (atomic number 8) is 1s2 2s2 2p4.