Introduction to Chemistry: 1st Semester Midterm Review Study Notes

Chapter 2 – Chemistry and Measurements

Metric and SI Units

Understanding the metric and SI (International System of Units) is fundamental in chemistry for accurate measurement and communication.

• Metric Prefixes: Common prefixes include kilo- (103), centi- (10-2), and milli- (10-3).

• SI Units: The base units include meter (m) for length, kilogram (kg) for mass, and second (s) for time.

Significant Figures

• Definition: Digits in a measurement that are known with certainty plus one estimated digit.

• Rules: Nonzero digits are always significant; zeros between nonzero digits are significant; leading zeros are not significant; trailing zeros are significant only if there is a decimal point.

• Example: 0.00450 has three significant figures.

Unit Conversions

• Dimensional Analysis: A method to convert one unit to another using conversion factors.

• Example: To convert 5.0 cm to meters: $5.0\,\text{cm} \times \frac{1\,\text{m}}{100\,\text{cm}} = 0.050\,\text{m}$

Density Calculations

• Formula: $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

• Application: Used to identify substances and solve for mass or volume.

Chapter 3 – Matter and Energy

Classification of Matter

• Pure Substances: Elements and compounds with fixed composition.

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

Physical vs. Chemical Properties

• Physical Properties: Observed without changing the substance (e.g., melting point, density).

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

Temperature Measurement

• Scales: Celsius (°C), Kelvin (K), Fahrenheit (°F).

• Conversion: $K = °C + 273.15$

Heat and Specific Heat

• Specific Heat: Amount of heat required to raise the temperature of 1 g of a substance by 1°C.

• Formula: $q = m \times c \times \Delta T$

• Example: Calculating heat absorbed by water when heated.

Chapter 4 – Elements and Symbols

Periodic Table Organization

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

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

Element Symbols

• Notation: One or two-letter abbreviations (e.g., H for hydrogen, Na for sodium).

Common Elements

• Examples: Hydrogen (H), Carbon (C), Oxygen (O), Nitrogen (N), Sodium (Na), Chlorine (Cl).

Atomic Structure

• Subatomic Particles: Protons (positive), neutrons (neutral), electrons (negative).

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

Chapter 5 – Electronic Structure of Atoms and Periodic Trends

Electron Arrangement

• Energy Levels: Electrons occupy specific energy levels (shells) around the nucleus.

• Electron Configuration: Distribution of electrons among orbitals (e.g., 1s2 2s2 2p6).

Periodic Trends

• Atomic Radius: Increases down a group, decreases across a period.

• Ionization Energy: Energy required to remove an electron; increases across a period, decreases down a group.

• Electron Affinity: Tendency of an atom to gain electrons.

Chapter 6 – Ionic and Molecular Compounds

Ionic Compounds

• Formation: Transfer of electrons from metals to nonmetals forms ions.

• Bond: Electrostatic attraction between cations and anions.

• Example: NaCl (sodium chloride)

Molecular Compounds

• Formation: Sharing of electrons between nonmetals.

• Bond: Covalent bond.

• Example: H2O (water)

Naming Compounds

• Ionic: Name cation first, then anion (e.g., sodium chloride).

• Molecular: Use prefixes to indicate number of atoms (e.g., carbon dioxide).

Chapter 7 – Chemical Quantities

Mole Concept

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

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

Percent Composition

• Formula: $\%\,\text{composition} = \frac{\text{mass of element}}{\text{total mass of compound}} \times 100$

Empirical and Molecular Formulas

• Empirical Formula: Simplest whole-number ratio of atoms in a compound.

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

Chapter 8 – Chemical Reactions

Balancing Chemical Equations

• Law of Conservation of Mass: Atoms are neither created nor destroyed in a chemical reaction.

• Steps: Write correct formulas, balance atoms one at a time, check work.

Types of Reactions

• Synthesis, Decomposition, Single Replacement, Double Replacement, Combustion

Predicting Products

• Use activity series and solubility rules to predict products and their states.

Chapter 9 – Chemical Quantities in Reactions

Stoichiometry

• Definition: Calculation of reactants and products in chemical reactions using balanced equations.

• Steps: Convert quantities to moles, use mole ratios, convert to desired units.

Limiting Reactant

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

Theoretical Yield

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

Chapter 16 – Nuclear Chemistry

Types of Radiation

• Alpha (α): Helium nucleus, low penetration.

• Beta (β): Electron, moderate penetration.

• Gamma (γ): High-energy photon, high penetration.

Balancing Nuclear Equations

• Conservation: Mass number and atomic number must be balanced on both sides.

Half-Life

• Definition: Time required for half of a radioactive sample to decay.

• Formula: $N = N_0 \left(\frac{1}{2}\right)^{t/t_{1/2}}$

Fission vs. Fusion

• Fission: Splitting of a heavy nucleus into lighter nuclei.

• Fusion: Combining of light nuclei to form a heavier nucleus.