1. Atomic Structure

Subatomic Particles

The atom is composed of three main subatomic particles: protons, neutrons, and electrons. Each has distinct properties and roles in atomic structure.

• Proton: Positively charged particle (+1), mass ≈ 1 amu, located in the nucleus.

• Neutron: Neutral particle (0 charge), mass ≈ 1 amu, located in the nucleus.

• Electron: Negatively charged particle (−1), mass ≈ 0 amu, located in electron cloud around nucleus.

• Atomic number (Z): Number of protons in the nucleus; defines the element.

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

• Isotopes: Atoms of the same element (same Z) with different numbers of neutrons (different A).

Quantum Numbers

Quantum numbers describe the properties and locations of electrons in atoms.

• n (principal quantum number): Indicates energy level (n = 1, 2, 3, ...).

• l (azimuthal quantum number): Describes orbital shape (l = 0 for s, 1 for p, 2 for d, 3 for f).

• ml (magnetic quantum number): Specifies orbital orientation (−l to +l).

• ms (spin quantum number): Electron spin (+½ or −½).

2. Periodic Trends

The periodic table displays recurring trends in element properties as a function of atomic number.

• Atomic radius: Decreases across a period (left to right), 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: Decreases across a period, increases down a group.

3. Bonding & Compounds

Ionic and Covalent Bonds

• Ionic bonds: Formed between metals and nonmetals via electron transfer.

• Covalent bonds: Formed between nonmetals via electron sharing.

• Polar bonds: Covalent bonds with a difference in electronegativity between atoms.

Lewis Structures

1. Count valence electrons.

2. Arrange atoms (least electronegative in center).

3. Bond then fill octets.

Molecular Geometry (VSEPR Theory)

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shapes based on electron pair repulsion.

• Linear: 180° bond angle

• Trigonal planar: 120°

• Tetrahedral: 109.5°

• Trigonal bipyramidal: 90°, 120°

• Octahedral: 90°

4. Stoichiometry

Stoichiometry involves quantitative relationships in chemical reactions.

• Mole: particles (Avogadro's number).

• Molar mass: g/mol, obtained from the periodic table.

• Conversions:

  • grams ↔ moles ↔ molecules/atoms

• % composition:

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

• Limiting reactant: The reactant that runs out first, limiting the amount of product formed.

5. Chemical Reactions

Types of Reactions

• Combination (Synthesis): Two or more substances form one product.

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

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

• Double replacement: Exchange of ions between two compounds.

• Combustion: Substance reacts with O2, producing heat and light.

Balancing Equations

• Ensure atoms on both sides are equal.

• Use the smallest whole-number coefficients.

6. Thermochemistry

Thermochemistry studies energy changes in chemical reactions, especially heat transfer.

• Heat (q):

• Enthalpy change (ΔH):

• Exothermic: (heat released)

• Endothermic: (heat absorbed)

7. Gases

Gas Laws

• Boyle's Law: (constant T, n)

• Charles's Law: (constant P, n)

• Avogadro's Law: (constant P, T)

• Ideal Gas Law: (R = 0.08206 L·atm/mol·K)

STP conditions: 1 atm, 273 K

8. Solutions

• Molarity (M):

• Dilution formula:

• Electrolytes: Substances that conduct electricity in solution (ions present).

• Solubility rules: Nitrates (NO3−) and alkali metal salts are always soluble.

9. Acids & Bases

• Arrhenius definition: Acid = H+ donor, Base = OH− donor.

• Brønsted-Lowry definition: Acid = proton donor, Base = proton acceptor.

• pH:

• pOH:

• pH + pOH = 14$ (at 25°C)

• Strong acids: HCl, HBr, HI, HNO3, H2SO4, HClO4

• Strong bases: Group 1 and 2 hydroxides

10. Equilibrium

• Equilibrium constant (K):

• K > 1: Products favored at equilibrium.

• K < 1: Reactants favored at equilibrium.

11. Kinetics

• Rate law:

• Activation energy (Ea): Minimum energy required for a reaction to occur; barrier to reaction.

• Catalyst: Lowers activation energy (Ea), increasing reaction rate.

12. Electrochemistry

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Anode: Site of oxidation.

• Cathode: Site of reduction.

• Cell potential (Ecell):