Grade 11 Chemistry Exam Review Notes – Comprehensive Study Guide

Ch. 2.4 & 2.5: The Structure of the Atom

Quantum Mechanical Model & Subshells

The quantum mechanical model describes the probability of finding an electron in a certain region of space around the nucleus, known as an orbital. Electrons are arranged in shells, subshells, and orbitals, each with specific energy levels and shapes.

• Orbital: A region in space where there is a high probability of finding an electron.

• Principal Quantum Number (n): Indicates the main energy level or shell (n = 1, 2, 3, ...).

• Subshells: Each shell contains one or more subshells (s, p, d, f), which differ in shape and energy.

• Electron Configuration: The arrangement of electrons in an atom's orbitals, following the Aufbau principle, Pauli exclusion principle, and Hund's rule.

Example: The electron configuration of oxygen (Z = 8) is 1s2 2s2 2p4.

Orbital Types and Shapes

• s-orbital: Spherical shape, 1 orientation.

• p-orbital: Dumbbell shape, 3 orientations (px, py, pz).

• d-orbital: Cloverleaf shape, 5 orientations.

• f-orbital: Complex shapes, 7 orientations.

Writing Electron Configurations

• Electrons fill orbitals in order of increasing energy (Aufbau principle).

• Each orbital holds a maximum of 2 electrons with opposite spins (Pauli exclusion principle).

• Electrons occupy degenerate orbitals singly before pairing (Hund's rule).

Shorthand (Noble Gas) Notation: Use the previous noble gas in brackets to represent core electrons. E.g., Na: [Ne] 3s1

Ch. 3.7: The Periodic Table of Elements – Periodic Trends

Atomic Radius

The atomic radius is the distance from the nucleus to the boundary within which electrons spend 90% of their time.

• Down a group: Atomic radius increases due to additional electron shells.

• Across a period: Atomic radius decreases due to increased nuclear charge pulling electrons closer.

Ionization Energy

• The energy required to remove an electron from a gaseous atom.

• Down a group: Ionization energy decreases (electrons are farther from the nucleus).

• Across a period: Ionization energy increases (greater nuclear charge).

Electron Affinity & Electronegativity

• Electron Affinity: The energy change when an electron is added to a neutral atom.

• Electronegativity: The ability of an atom to attract electrons in a chemical bond.

• Both generally increase across a period and decrease down a group.

Ch. 3.8: Empirical and Molecular Formulas

Percentage Composition

• Gives the percentage by mass of each element in a compound.

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

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.

• To determine empirical formula: Convert mass to moles, divide by smallest number of moles, multiply to get whole numbers if necessary.

Example: A compound with 40% C, 6.7% H, and 53.3% O by mass has the empirical formula CH2O.

Ch. 4.1: Solutions and Solubility

Solutions

• Solution: Homogeneous mixture of solute dissolved in solvent.

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

Properties of Solutions

• Electrolytes: Substances that conduct electricity when dissolved in water (e.g., NaCl).

• Nonelectrolytes: Substances that do not conduct electricity in solution (e.g., sugar).

Intermolecular Forces

• London Dispersion Forces (LDF): Weak, present in all molecules.

• Dipole-Dipole Forces: Between polar molecules.

• Hydrogen Bonding: Strong dipole-dipole interaction involving H bonded to N, O, or F.

Like Dissolves Like: Polar solvents dissolve polar solutes; non-polar solvents dissolve non-polar solutes.

Ch. 4.2: Balancing Reactions & Types of Chemical Reactions

Balancing Chemical Equations

• Ensure the same number of each atom on both sides of the equation.

• Balance metals, then nonmetals, then hydrogen and oxygen last.

Types of Reactions

• Combustion: Hydrocarbon reacts with O2 to form CO2 and H2O.

• Synthesis: Two or more substances combine to form one product.

• Decomposition: One substance breaks down into two or more products.

• Single Displacement: One element replaces another in a compound.

• Double Displacement: Exchange of ions between two compounds.

Ch. 4.4.1: Double Displacement, Precipitation, and Acid-Base Reactions

Precipitation Reactions

• Occur when two aqueous solutions form an insoluble product (precipitate).

• Use solubility rules to predict precipitate formation.

Acid-Base Reactions

• Arrhenius Theory: Acids produce H+ in water, bases produce OH-.

• Brønsted-Lowry Theory: Acids donate protons, bases accept protons.

• General equation: Acid + Base → Salt + Water

Ch. 4.4.2: Oxidation-Reduction (Redox) Reactions

Redox Reactions

• Involve the transfer of electrons between species.

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Oxidation and reduction always occur together.

Example: Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s)

Ch. 4.4.3: Limiting Reagent, Theoretical Yield, and Percent Yield

Limiting Reagent

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

Theoretical and Actual Yield

• Theoretical Yield: Maximum amount of product possible from given reactants.

• Actual Yield: Amount of product actually obtained.

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

Ch. 4.4.4: Solution Concentration

Concentration Units

• Molarity (M): $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Percent by mass, volume, and parts per million (ppm): Used for very dilute solutions.

Preparing Solutions

• Dissolving a solid solute in a liquid solvent or diluting a concentrated solution.

• Use volumetric flasks and pipettes for precise measurements.

Ch. 4.5: Solution Stoichiometry and Acid-Base Titration

Solution Stoichiometry

• Relates the volumes and concentrations of solutions to the amount of reactants and products.

• Key formula: $n = C \times V$ (where n = moles, C = concentration, V = volume in L)

Acid-Base Titration

• Used to determine the concentration of an acid or base by reacting it with a standard solution of known concentration.

• At the equivalence point, moles of acid = moles of base.

Ch. 12.6: Colligative Properties

Colligative Properties

• Depend on the number of solute particles, not their identity.

• Include vapor pressure lowering, boiling point elevation, freezing point depression, and osmotic pressure.

• Raoult's Law: $P_{\text{solution}} = X_{\text{solvent}} P^0_{\text{solvent}}$

Ch. 13.1: The Simple Gas Laws

Gas Laws

• Boyle's Law: $P_1V_1 = P_2V_2$ (at constant T, n)

• Charles' Law: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$ (at constant P, n)

• Combined Gas Law: $\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$

Mixtures of Gases and Partial Pressures

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

Ch. 3.1: Organic Chemistry – Hydrocarbons and Functional Groups

Hydrocarbons

• Alkanes: Single bonds only, general formula CnH2n+2

• Alkenes: At least one double bond, general formula CnH2n

• Alkynes: At least one triple bond, general formula CnH2n-2

Functional Groups

• Alcohols: Contain -OH group

• Halides: Contain halogen atoms (F, Cl, Br, I)

• Aldehydes and Ketones: Contain carbonyl group (C=O)

• Carboxylic Acids: Contain -COOH group

• Amines: Contain -NH2 group

Ch. 11.1: Phase Diagrams

Phase Diagrams

• Show the state of a substance (solid, liquid, gas) at different temperatures and pressures.

• Triple Point: All three phases coexist in equilibrium.

• Critical Point: The highest temperature and pressure at which a liquid can exist.

Ch. 11.2: Crystalline Solids – Fundamental Types

Types of Crystalline Solids

• Ionic: Composed of ions, high melting points, conduct electricity when molten.

• Molecular: Composed of molecules, low melting points, do not conduct electricity.

• Covalent Network: Atoms connected by covalent bonds, very high melting points.

• Metallic: Metal atoms, good conductors, malleable and ductile.

Ch. 12.1: Energetics of Solution Formation

Enthalpy of Solution

• Formation of a solution involves breaking solute and solvent interactions (endothermic) and forming new solute-solvent interactions (exothermic).

• Enthalpy of Solution ($\Delta H_{\text{soln}}$): $\Delta H_{\text{soln}} = \Delta H_1 + \Delta H_2 + \Delta H_3$

Ch. 13.2: Kinetic Molecular Theory

Kinetic Molecular Theory (KMT)

• Gases consist of small particles in constant, random motion.

• Collisions between particles are elastic (no energy loss).

• Volume of gas particles is negligible compared to the container.

• Average kinetic energy is proportional to temperature.

Summary Table Example: Types of Intermolecular Forces

Type	Strength	Occurs Between
London Dispersion	Weakest	All molecules
Dipole-Dipole	Intermediate	Polar molecules
Hydrogen Bonding	Strongest (of IMFs)	H bonded to N, O, or F

Additional info: These notes cover key concepts from atomic structure, periodic trends, chemical reactions, stoichiometry, solutions, gas laws, organic chemistry, and intermolecular forces, providing a comprehensive review for a General Chemistry course.