General Chemistry Final Exam Study Guide

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This section covers foundational concepts in chemistry, including definitions, classifications, and basic properties of matter.

Matter: Anything that has mass and occupies space.
Atoms, Elements, Compounds: Atoms are the basic units of matter; elements are pure substances consisting of one type of atom; compounds are substances formed from two or more elements chemically bonded.
States of Matter: Solid, liquid, and gas; differences include shape, volume, and particle arrangement.
Mixtures vs. Pure Substances: Mixtures contain two or more substances physically combined; pure substances have a fixed composition.
Homogeneous vs. Heterogeneous Mixtures: Homogeneous mixtures have uniform composition; heterogeneous mixtures do not.
Physical vs. Chemical Properties: Physical properties can be observed without changing the substance; chemical properties describe a substance's ability to undergo chemical changes.

Understanding measurement systems and conversions is essential for quantitative chemistry.

SI Units: Standard units for length (meter), mass (kilogram), time (second), temperature (kelvin), amount of substance (mole).
Unit Conversion: Use conversion factors to change between units.
Temperature Scales: Celsius, Kelvin, Fahrenheit; conversion formulas:
Density:
Exact vs. Inexact Numbers: Exact numbers are counted or defined; inexact numbers are measured.

Proper reporting of measurements is crucial for scientific communication.

Accuracy: Closeness to the true value.
Precision: Reproducibility of measurements.
Significant Figures: Digits that carry meaning in a measurement; rules for addition, subtraction, multiplication, and division.

Dimensional analysis is used to convert units and solve quantitative problems.

Conversion Factors: Ratios used to express the same quantity in different units.
Multiple Step Conversions: Chain conversion factors to solve complex problems.

Atomic theory explains the nature and behavior of atoms.

Dalton's Atomic Theory: Matter is composed of atoms; atoms of each element are identical; atoms combine in simple ratios to form compounds.
Law of Conservation of Mass: Mass is conserved in chemical reactions.
Law of Definite Proportions: Compounds have fixed ratios of elements.
Law of Multiple Proportions: Elements can combine in different ratios to form different compounds.
Subatomic Particles: Protons (+), neutrons (0), electrons (-); location and properties.
Isotopes: Atoms of the same element with different numbers of neutrons.
Atomic Number (): Number of protons.
Mass Number (): Number of protons plus neutrons.
Average Atomic Mass: Weighted average based on isotopic abundance.

The periodic table organizes elements by atomic number and properties.

Groups and Periods: Groups are columns; periods are rows.
Metals, Nonmetals, Metalloids: Location and properties.
Special Group Names: Alkali metals (Group 1), alkaline earth metals (Group 2), halogens (Group 17), noble gases (Group 18).
Periodic Trends: Atomic radius, ionization energy, electron affinity, electronegativity.
Effective Nuclear Charge (): (where is the shielding constant).

Chemical bonds form between atoms to create molecules and compounds.

Ionic Bonds: Formed by transfer of electrons between metals and nonmetals.
Covalent Bonds: Formed by sharing electrons between nonmetals.
Polar vs. Nonpolar Covalent Bonds: Determined by electronegativity difference.
Electronegativity: Tendency of an atom to attract electrons in a bond.
Lewis Dot Structures: Visual representation of valence electrons and bonding.
Formal Charge:
Resonance Structures: Multiple valid Lewis structures for a molecule.
VSEPR Theory: Predicts molecular geometry based on electron pair repulsion.
Hybridization: Mixing of atomic orbitals to form new hybrid orbitals.
Bond Strength and Length: Multiple bonds are stronger and shorter than single bonds.

Stoichiometry involves quantitative relationships in chemical reactions.

Balancing Chemical Equations: Ensure the same number of atoms of each element on both sides.
Types of Reactions: Synthesis, decomposition, single replacement, double replacement, combustion.
Limiting Reactant: The reactant that is completely consumed first.
Theoretical Yield: Maximum amount of product possible.
Percent Yield:
Molarity:
Solution Preparation: Calculating concentrations and dilutions.

Solutions can conduct electricity depending on the presence of ions.

Electrolytes: Substances that dissociate into ions in solution; strong, weak, and nonelectrolytes.
Precipitation Reactions: Formation of an insoluble product.
Acid-Base Reactions: Neutralization between acids and bases.
Redox Reactions: Transfer of electrons; oxidation and reduction.
Displacement Reactions: Single and double displacement.
Calculating Molarity: Use mass of solute and volume of solution.

Gases have unique properties and behaviors described by gas laws.

Quantum mechanics explains the behavior of electrons in atoms.

Wavelength () and Frequency ():
Electromagnetic Spectrum: Range of all types of electromagnetic radiation.
Energy of a Photon:
Quantum Numbers: Principal (), angular (), magnetic (), spin ().
Electron Configuration: Arrangement of electrons in orbitals; use Aufbau principle, Pauli exclusion, Hund's rule.
Periodic Trends Explained by Quantum Theory: Atomic radius, ionization energy, electron affinity.

Molecular geometry and bonding theories predict the shapes and properties of molecules.

This guide covers all major topics listed in a standard General Chemistry curriculum, including measurement, atomic theory, periodic trends, chemical bonding, stoichiometry, solutions, gases, and quantum mechanics.
Students should be familiar with definitions, formulas, and problem-solving strategies for each topic.