CHEM 1025 Final Exam Review Topics (Ch 2-10 & 13): Structured Study Notes

Chapter 2: Measurement and Problem Solving

Scientific Measurement

Accurate measurement is fundamental in chemistry for quantifying substances and reactions.

• Significant Figures: Digits in a measurement that are known with certainty plus one estimated digit. Used to express precision.

• Basic Units of Measurement: Standard units include meter (length), kilogram (mass), second (time), mole (amount of substance).

• Unit Conversions: Changing from one unit to another using conversion factors.

• Density: The ratio of mass to volume, calculated as $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$.

• Scientific Notation: Expressing numbers as a product of a coefficient and a power of ten, e.g., $3.0 \times 10^4$.

Chapter 3: Matter and Energy

Matter Classification

Matter is anything that has mass and occupies space. It can be classified based on composition and properties.

• Pure Substances: Elements and compounds with uniform composition.

• Mixtures: Physical combinations of substances; can be homogeneous (uniform) or heterogeneous (non-uniform).

• Physical vs. Chemical Properties: Physical properties can be observed without changing the substance; chemical properties involve changes in composition.

• Energy: The capacity to do work; includes kinetic and potential energy.

Chapter 4: Atoms and Elements

Atomic Structure and Models

Atoms are the basic units of matter, composed of protons, neutrons, and electrons.

• Dalton’s Atomic Theory: Matter is made of indivisible atoms; atoms of each element are identical.

• Thomson and the Plum Pudding Model: Electrons are embedded in a positively charged sphere.

• Rutherford’s Gold Foil Experiment: Discovered the nucleus as a dense, positively charged center.

• Periodic Table: Organizes elements by increasing atomic number; groups (columns) and periods (rows) indicate similar properties.

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

Chapter 5: Molecules and Compounds

Chemical Bonding and Compound Types

Atoms combine to form compounds through chemical bonds.

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

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

• Chemical Formulas: Represent the types and numbers of atoms in a compound, e.g., $\text{H}_2\text{O}$.

Chapter 6: Chemical Composition

Quantifying Chemical Substances

Chemists use the mole to count particles and relate mass to number of atoms or molecules.

• Mole Concept: 1 mole = $6.022 \times 10^{23}$ particles (Avogadro’s number).

• Molar Mass: Mass of one mole of a substance, calculated from the periodic table.

• Percent Composition: Percentage by mass of each element in a compound.

• Empirical and Molecular Formulas: Empirical formula shows simplest ratio; molecular formula shows actual number of atoms.

Chapter 7: Chemical Reactions

Types and Representation of Chemical Reactions

Chemical reactions involve the transformation of substances into new products.

• Balancing Equations: Ensures the same number of each atom on both sides of the equation.

• Types of Reactions: Synthesis, decomposition, single displacement, double displacement, combustion.

• Precipitation Reactions: Formation of an insoluble product from soluble reactants.

• Acid-Base Reactions: Transfer of protons between substances.

• Redox Reactions: Transfer of electrons; oxidation (loss), reduction (gain).

Chapter 8: Quantities in Chemical Reactions

Stoichiometry

Stoichiometry involves calculating the amounts of reactants and products in chemical reactions.

• Mole Ratios: Derived from balanced equations to relate quantities of substances.

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

• Theoretical Yield: Maximum amount of product possible.

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

Chapter 9: Electrons in Atoms and the Periodic Table

Atomic Structure and Electron Configuration

Electron arrangement determines chemical properties and periodic trends.

• Orbitals: Regions where electrons are likely to be found; s, p, d, f types.

• Electron Configuration: Distribution of electrons among orbitals, e.g., $1s^2 2s^2 2p^6$.

• Periodic Trends: Patterns in properties such as atomic radius, ionization energy, and electronegativity.

Chapter 10: Chemical Bonding

Bond Types and Molecular Geometry

Chemical bonds hold atoms together in compounds.

• Ionic Bonds: Electrostatic attraction between oppositely charged ions.

• Covalent Bonds: Sharing of electron pairs between atoms.

• Lewis Structures: Diagrams showing bonding and lone pairs.

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

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

• Polarity: Unequal sharing of electrons leads to polar molecules.

Chapter 13: Solutions

Properties and Calculations Involving Solutions

Solutions are homogeneous mixtures of solute and solvent.

• Solubility: Maximum amount of solute that can dissolve in a solvent at a given temperature.

• Concentration Units: Includes molarity ($M = \frac{\text{moles of solute}}{\text{liters of solution}}$), molality, and percent composition.

• Solution Preparation: Calculating amounts needed to prepare solutions of desired concentration.

• Types of Solutions: Saturated, unsaturated, and supersaturated.

Periodic Table Reference

Organization and Use

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

• Groups: Vertical columns; elements in the same group have similar chemical properties.

• Periods: Horizontal rows; properties change progressively across a period.

• Metals, Nonmetals, Metalloids: Classified based on physical and chemical properties.

Classification	Properties
Metals	Conduct electricity, malleable, shiny
Nonmetals	Poor conductors, brittle, dull
Metalloids	Intermediate properties