General Chemistry I: Midterm Study Guide (Chapters 1-5)

Chapter 1: Units, Measurement & Matter

Units of Measurement

Understanding units is fundamental in chemistry for quantifying physical properties and performing calculations.

• SI Units: The International System of Units (SI) is the standard for scientific measurements. Common SI units include meter (m) for length, kilogram (kg) for mass, and second (s) for time.

• Metric & U.S. Conversions: Converting between metric and U.S. customary units is essential for interpreting data and solving problems.

• Dimensional Analysis: A method for converting between units using conversion factors.

Example: To convert 10 inches to centimeters, use the conversion factor: $1\ \text{in} = 2.54\ \text{cm}$.

Significant Figures & Scientific Notation

Significant figures reflect the precision of a measurement, while scientific notation expresses very large or small numbers efficiently.

• Identifying significant figures: All nonzero digits are significant; zeros between nonzero digits are significant; leading zeros are not significant.

• Scientific notation: Numbers are written as $a \times 10^n$, where $1 \leq a < 10$ and $n$ is an integer.

Example: 0.00456 has three significant figures and can be written as $4.56 \times 10^{-3}$.

Properties of Matter

Matter is anything that has mass and occupies space. Its properties can be classified as physical or chemical.

• Physical properties: Characteristics observed without changing the substance (e.g., color, melting point).

• Chemical properties: Characteristics observed during a chemical change (e.g., flammability).

Classification of Matter

Matter can be classified based on its composition.

• Pure substances: Elements and compounds with uniform composition.

• Mixtures: Combinations of two or more substances (homogeneous or heterogeneous).

Chapter 2: Atoms, Elements & The Periodic Table

Atomic Structure

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

• Protons: Positively charged particles in the nucleus.

• Neutrons: Neutral particles in the nucleus.

• Electrons: Negatively charged particles orbiting the nucleus.

Example: Carbon-12 has 6 protons, 6 neutrons, and 6 electrons.

The Periodic Table

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

• Groups: Vertical columns with similar chemical properties.

• Periods: Horizontal rows.

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

Electron Configuration

Electron configuration describes the arrangement of electrons in an atom.

• Valence electrons: Electrons in the outermost shell, important for chemical bonding.

• Example: Sodium (Na): $1s^2 2s^2 2p^6 3s^1$

Chapter 3: Bonding, Naming & Molecular Geometry

Ionic & Covalent Bonds

Chemical bonds form when atoms share or transfer electrons.

• Ionic bonds: Formed by transfer of electrons from metals to nonmetals.

• Covalent bonds: Formed by sharing electrons between nonmetals.

Naming Compounds

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

• Covalent compounds: Use prefixes to indicate the number of atoms (e.g., CO2: carbon dioxide).

Lewis Structures & VSEPR Geometry

Lewis structures represent valence electrons and predict molecular geometry using the Valence Shell Electron Pair Repulsion (VSEPR) theory.

• VSEPR theory: Electron pairs around a central atom arrange themselves to minimize repulsion, determining molecular shape.

Example: Water (H2O) has a bent geometry due to two lone pairs on oxygen.

Chapter 4: Chemical Reactions & Stoichiometry

Writing & Interpreting Formulas

Chemical formulas represent the elements and their ratios in compounds.

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

• Molecular formula: Actual number of atoms of each element.

Balancing Chemical Equations

Balancing ensures the same number of atoms of each element on both sides of the equation.

• Law of Conservation of Mass: Matter is neither created nor destroyed in a chemical reaction.

Example: $2H_2 + O_2 \rightarrow 2H_2O$

Stoichiometry

Stoichiometry involves quantitative relationships between reactants and products.

• Mole concept: 1 mole = $6.022 \times 10^{23}$ particles.

• Molar mass: Mass of 1 mole of a substance (g/mol).

• Mole-mole, mass-mass, and mole-mass conversions: Used to calculate amounts in reactions.

Example: To find grams of water produced from 4 moles of hydrogen: $4\ \text{mol} \times \frac{2\ \text{mol}\ H_2O}{2\ \text{mol}\ H_2} \times 18.02\ \text{g/mol} = 72.08\ \text{g}$

Chapter 5: States of Matter & Gas Laws

Properties of Solids, Liquids, Gases

Matter exists in three primary states, each with distinct properties.

• Solids: Definite shape and volume; particles closely packed.

• Liquids: Definite volume, indefinite shape; particles less tightly packed.

• Gases: Indefinite shape and volume; particles far apart and move freely.

Gas Laws

Gas laws describe the relationships between pressure, volume, temperature, and amount of gas.

• 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)

• Avogadro's Law: $\frac{V_1}{n_1} = \frac{V_2}{n_2}$ (at constant temperature and pressure)

• Ideal Gas Law: $PV = nRT$

• Dalton's Law of Partial Pressures: $P_{\text{total}} = P_1 + P_2 + P_3 + \ldots$

Example: Calculate the pressure exerted by 2 moles of gas in a 5 L container at 300 K: $P = \frac{nRT}{V}$

Crystalline Solids

Crystalline solids have an ordered arrangement of particles, leading to distinct physical properties.

• Types: Ionic, molecular, covalent network, metallic.

Gas Pressure Units

Pressure can be measured in several units, including atmospheres (atm), pascals (Pa), and millimeters of mercury (mmHg).

• Conversions: $1\ \text{atm} = 101,325\ \text{Pa} = 760\ \text{mmHg}$

State of Matter	Shape	Volume	Particle Arrangement
Solid	Definite	Definite	Closely packed, orderly
Liquid	Indefinite	Definite	Close, but not fixed
Gas	Indefinite	Indefinite	Far apart, random