General Chemistry I: Comprehensive Study Notes (CHEM 120 Final Exam Practice Questions)

Atomic Structure

Subatomic Particles and Atomic Number

The atom is composed of three fundamental particles: protons, neutrons, and electrons. The atomic number (Z) is the number of protons in the nucleus and defines the element.

• Protons: Positively charged particles found in the nucleus.

• Neutrons: Neutral particles found in the nucleus.

• Electrons: Negatively charged particles found in orbitals around the nucleus.

• In a neutral atom, the number of protons equals the number of electrons.

• Mass number (A): Total number of protons and neutrons in the nucleus.

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

Isotopes and Ions

Isotopes are atoms of the same element with different numbers of neutrons. Ions are atoms or molecules that have gained or lost electrons.

• Cation: Positively charged ion (loss of electrons).

• Anion: Negatively charged ion (gain of electrons).

• Isotopes have the same atomic number but different mass numbers.

Example: $^{35}Cl$ and $^{37}Cl$ are isotopes of chlorine.

Electronic Configuration

Electrons are arranged in shells and subshells according to the Aufbau principle, Pauli exclusion principle, and Hund's rule.

• Aufbau Principle: Electrons fill the lowest energy orbitals first.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers.

• Hund's Rule: Electrons occupy orbitals singly before pairing.

• Electron configuration notation: $1s^2 2s^2 2p^6 3s^2 3p^6$ (for Argon).

Example: The ground-state electron configuration of Mn is $1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^5$.

Molecular Structure and Bonding

Lewis Structures and Molecular Geometry

Lewis structures represent the arrangement of electrons in molecules. Molecular geometry is determined by the number of bonding and lone pairs around the central atom (VSEPR theory).

• VSEPR Theory: Valence Shell Electron Pair Repulsion theory predicts molecular shapes.

• Common geometries: linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.

• Bond angles depend on geometry and lone pairs.

Example: $BF_3$ is trigonal planar; $NH_3$ is trigonal pyramidal.

Types of Chemical Bonds

Chemical bonds include ionic, covalent, and metallic bonds.

• Ionic bond: Transfer of electrons from metal to nonmetal.

• Covalent bond: Sharing of electrons between nonmetals.

• Polar covalent bond: Unequal sharing of electrons.

• Hybridization: Mixing of atomic orbitals to form new hybrid orbitals (e.g., $sp^3$, $sp^2$, $sp$).

Example: $CO_2$ is linear and has double bonds between C and O.

Stoichiometry

Mole Concept and Molar Mass

The mole is the SI unit for amount of substance. Molar mass is the mass of one mole of a substance.

• Avogadro's number: $6.022 \times 10^{23}$ particles/mol.

• Molar mass (g/mol) is used to convert between mass and moles.

Example: 1 mole of $H_2O$ has a mass of 18.02 g.

Empirical and Molecular Formulas

Empirical formula shows the simplest whole-number ratio of atoms. Molecular formula shows the actual number of atoms in a molecule.

• Determine empirical formula from percent composition.

• Molecular formula = (empirical formula) × n, where n is a whole number.

Example: $C_2H_6$ (ethane) has the empirical formula $CH_3$.

Stoichiometric Calculations

Stoichiometry involves calculations based on balanced chemical equations.

• Use coefficients to relate moles of reactants and products.

• Limiting reagent: The reactant that is completely consumed first.

• Percent yield: $\frac{\text{actual yield}}{\text{theoretical yield}} \times 100\%$

Example: In $2H_2 + O_2 \rightarrow 2H_2O$, 2 moles of $H_2$ react with 1 mole of $O_2$ to produce 2 moles of $H_2O$.

States of Matter and Solutions

Properties of Gases, Liquids, and Solids

Matter exists as solids, liquids, and gases. Each state has distinct properties.

• Gases: No fixed shape or volume, compressible, particles far apart.

• Liquids: Fixed volume, no fixed shape, particles closer together.

• Solids: Fixed shape and volume, particles tightly packed.

Example: Water vapor, liquid water, and ice are the three states of $H_2O$.

Solution Concentration

Molarity (M) is the most common unit of concentration, defined as moles of solute per liter of solution.

• $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Preparation of solutions involves dissolving a known mass of solute and diluting to a specific volume.

Example: Dissolving 5.84 g of NaCl in water to make 1.00 L solution gives a 0.100 M NaCl solution.

Energetics (Thermochemistry)

Heat, Work, and Internal Energy

Thermochemistry studies energy changes in chemical reactions. Heat (q) and work (w) are forms of energy transfer.

• First Law of Thermodynamics: $\Delta E = q + w$

• Enthalpy (H): $\Delta H = \Delta E + P\Delta V$

• Endothermic reactions absorb heat; exothermic reactions release heat.

Example: Combustion of methane is exothermic.

Calorimetry and Specific Heat

Calorimetry measures heat changes. Specific heat (c) is the amount of heat required to raise the temperature of 1 g of a substance by 1°C.

• $q = mc\Delta T$

• Use calorimeters to determine heat of reaction.

Example: Heating 100 g of water by 10°C requires $q = 100 \times 4.18 \times 10 = 4180$ J.

Standard Enthalpy Changes

Standard enthalpy of formation ($\Delta H_f^\circ$) is the enthalpy change when 1 mole of a compound forms from its elements in their standard states.

• Hess's Law: The total enthalpy change is the sum of enthalpy changes for individual steps.

• Bond enthalpy: Energy required to break a bond in a molecule.

Example: $\Delta H_{rxn} = \sum \Delta H_{f,\text{products}}^\circ - \sum \Delta H_{f,\text{reactants}}^\circ$

Descriptive Chemistry and Periodic Properties

Periodic Table Trends

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

• Atomic radius decreases across a period, increases down a group.

• Ionization energy increases across a period, decreases down a group.

• Electronegativity increases across a period, decreases down a group.

• Metallic character increases down a group, decreases across a period.

Example: Fluorine is the most electronegative element.

Reactivity and Chemical Properties

Elements show characteristic reactivity based on their position in the periodic table.

• Alkali metals are highly reactive with water.

• Halogens are reactive nonmetals.

• Noble gases are largely inert.

Example: Sodium reacts vigorously with water to produce $NaOH$ and $H_2$ gas.

Laboratory Skills

Measurement and Precision

Accurate measurement is essential in chemistry. Precision refers to the reproducibility of measurements; accuracy refers to closeness to the true value.

• Use balances, graduated cylinders, and volumetric flasks for measurement.

• Significant figures indicate the precision of a measurement.

Example: Mass of water measured as 1.298 g (3 significant figures).

Law of Conservation of Mass

The Law of Conservation of Mass states that mass is neither created nor destroyed in a chemical reaction.

• Mass of reactants equals mass of products.

• Used to verify experimental results.

Example: Burning Mg ribbon in O2 and comparing mass before and after reaction.

Tables

Sample Table: Molar Masses of Compounds

Formula	Molar Mass (g/mol)
CaCO3	100.1
Na2SO4	142.1
NaCl	58.5
Na2CO3	106.0

Sample Table: Ion Radii

Ion	Ionic Radius (nm)
Na+	0.095
Mg2+	0.065
Al3+	0.050

Sample Table: Standard Enthalpy of Formation

Compound	ΔHf° (kJ/mol)
CO2(g)	-393.5
H2O(l)	-285.8
CH4(g)	-74.8

Additional info: These study notes are based on a comprehensive set of final exam practice questions covering core topics in General Chemistry I, including atomic structure, molecular bonding, stoichiometry, states of matter, thermochemistry, periodic properties, and laboratory techniques. All equations are provided in LaTeX format for clarity and academic rigor.