1.) Combination: A+B-AB

2.) Decomposition: AB-A+B

3.) Single Displacement: Ax+B-A+Bx

4.) Double Displacement: Ax+By-Ay+Bx

5.) Combustion: AB+O-AO+BO

Things to Remember:

1.) Determine the formula weight

Unit: Atomic Mass Unit(amu)

2.) Interconverting Masses, Moles, and number of entities

Grams to Moles and vice versa: Use molar mass (Bridge 1)

Moles to Formula units and vice versa: Use Avogadro's number (Bridge 2)

1 mol= 6.02 x 10^23 entities

Postulates:

1.) cannot be created nor destroyed

2.) one element cannot be converted into another element

3.) identical in mass and other properties(unique)

4.) Chemical combination of a specific ratio of atom of different elements form compound

ATOMIC THEORY AND MASS LAWS

1.) Law of Conservation of Mass- Postulates 1, 2, and 3

2.) Law of Definite Composition- Postulates 3 and 4

3.) Law of Multiple Proportions- Postulates 1, 3, and 4

Chemical formula- number of atoms of each element in a compound or molecule

Empirical formula- simplest whole number ratio of elements

Molecular formula- actual whole number ratio

WRITING FORMULA OF IONIC COMPOUND

1.) Determine Cation and charge

2.) Determine Anion and charge

3.) Do crisscross method

4.) Simplify if necessary

NAMING ACIDS

1.) If the anion ends in -ide:

• change -ide to -ic

• Add hydro-

• Add the word "acid"

2.) If the anion of the acid ends in -ate or -ite

• change -ate to -ic or -ite to -ous

• add the word "acid"

NAMING CATIONS

1.) Cations formed from metal atoms have the same name as the metal

2.) If a metal can form cations with different charges, the positive charge is indicated by a Roman Numeral in parenthesis following the name of the metal.

NAMING ANIONS

1.) ending: -ide

2.) -ate: most common or representative oxyanion

-ite: one O atom fewer

2.1) perchlorate, chlorate, chlorite, hypochlorite

NAMING IONIC COMPOUNDS

1.) cation name (space) followed by the anion name

Mono - 1 Hexa - 6

Di - 2 Hepta - 7

Tri - 3 Octa - 8

Tetra - 4 Nona - 9

Penta - 5 Deca - 10

NAMING BINARY MOLECULAR COMPOUNDS

1. Closest to metals is written first

2. Second element is given -ide ending

3. mono- is never used with first element

4. Drop o/a if repeated

5. lower one is named first if same group

• John Dalton (1803): Proposed that all matter is made of atoms.

• Postulates:

  1. Elements are made of indivisible atoms.

  2. Atoms of the same element are identical.

  3. Atoms of different elements differ in properties.

  4. Compounds are formed by combinations of atoms in whole-number ratios.

  5. Chemical reactions involve rearrangement of atoms, not their creation or destruction.

• Limitations:

  • Atoms are divisible (protons, neutrons, electrons).

  • Isotopes exist (atoms of the same element with different masses).

2. Quantum Numbers

• Describe electron positions and behavior in atoms.

• Principal (n): Energy level (1, 2, 3…).

• Angular Momentum (l): Shape of orbital (s = 0, p = 1, d = 2, f = 3).

• Magnetic (ml): Orbital orientation (−l to +l).

• Spin (ms): Spin of electron (+½ or −½).

3. Periodic Table

• Arranged by increasing atomic number (Moseley).

• Groups: Columns with similar properties.

• Periods: Rows indicating energy levels.

• Trends:

  • Atomic radius: Increases down, decreases across.

  • Ionization energy & Electronegativity: Decrease down, increase across.

  • Metallic character: Increases down, decreases across.

4. Types of Chemical Bonds

• Ionic: Transfer of electrons (metal + nonmetal).

• Covalent: Sharing electrons (nonmetal + nonmetal).

• Metallic: Delocalized electrons among metal atoms.

• Polar vs. Nonpolar Covalent:

  • Polar: Unequal sharing.

  • Nonpolar: Equal sharing.

5. Chemical Structures

• Lewis Structures: Show bonding and lone pairs.

• Octet Rule: Atoms tend to gain 8 valence electrons (exceptions exist).

• Use dots for lone pairs, lines for shared pairs.

6. Valence Bond Theory

• Bonds form when orbitals of two atoms overlap.

• Hybridization: Mixing of orbitals (e.g., sp³ for tetrahedral).

• Explains bond angles and molecular shapes.

7. Molecular Geometry (VSEPR Theory)

• Shape based on repulsion between electron pairs.

• Common Geometries:

  • Linear (2 pairs)

  • Trigonal Planar (3 pairs)

  • Tetrahedral (4 pairs)

  • Trigonal Bipyramidal (5 pairs)

  • Octahedral (6 pairs)

8. Resonance

• Occurs when more than one valid Lewis structure exists.

• Electrons are delocalized.

• Example: O₃, CO₃²⁻

9. Kinetic Molecular Theory (KMT)

• Gases consist of particles in constant, random motion.

• Collisions are elastic (no energy loss).

• Gas particles occupy no volume and exert no forces (ideal gas).

• Temperature is proportional to average kinetic energy.

10. Gas Laws

• Boyle’s Law: P₁V₁ = P₂V₂ (constant T)

• Charles’s Law: V₁/T₁ = V₂/T₂ (constant P)

• Avogadro’s Law: V ∝ n (constant T, P)

• Combined Gas Law: (P₁V₁)/T₁ = (P₂V₂)/T₂

11. Ideal Gas Law & Dalton’s Law

• PV = nRT

  • P = pressure, V = volume, n = moles, R = 0.0821 L·atm/mol·K, T = temperature

• Dalton’s Law of Partial Pressures:

  • P_total = P₁ + P₂ + … + Pₙ

  • P_gas = (mol_gas/total mol) × P_total

12. Graham’s Law of Diffusion

• Rate₁/Rate₂ = √(M₂/M₁)

  • M = molar mass

  • Lighter gases diffuse faster.

13. Gay-Lussac’s Law

• P₁/T₁ = P₂/T₂ (constant V)

• Pressure of gas is directly proportional to temperature.

14. Gas Stoichiometry

• Use Ideal Gas Law or molar volume (22.4 L at STP) to convert between gas volume and moles.

• Combine with mole ratios from balanced equations.

• Forces between molecules, weaker than chemical bonds.

• Types:

  • London Dispersion Forces: Present in all molecules, dominant in nonpolar.

  • Dipole-Dipole Forces: Between polar molecules.

  • Hydrogen Bonding: Strong dipole-dipole interaction between H and N, O, or F.

  • Ion-Dipole: Between ion and polar molecule (common in solutions).

2. Physical Properties of Liquids Affected by IMFA

• Boiling Point: Stronger IMFA → higher boiling point.

• Viscosity: Resistance to flow; increases with stronger IMFA.

• Surface Tension: Tendency to minimize surface area; stronger with higher IMFA.

• Vapor Pressure: Decreases with stronger IMFA.

• Capillary Action: Movement in narrow tubes; due to cohesion and adhesion.

3. Structure and Properties of Water

• Polar molecule, bent shape (~104.5°).

• Forms hydrogen bonds.

• High boiling point, heat capacity, and surface tension.

• Universal solvent due to polarity.

• Solid (ice) is less dense than liquid due to open hexagonal lattice.

4. Types and Properties of Solids

• Crystalline: Ordered, repeating structure (e.g., NaCl, quartz).

  • Types:

    • Ionic: High melting point, brittle, conducts when molten.

    • Covalent Network: Very hard, high melting point (e.g., diamond).

    • Metallic: Malleable, conductive, delocalized electrons.

    • Molecular: Soft, low melting point (e.g., sugar).

• Amorphous: No long-range order (e.g., glass, plastic).

5. Important Features of Phase Diagram

• Phase diagram: Graph showing states of matter under various conditions.

  • Axes: Pressure (y-axis) vs. Temperature (x-axis).

• Lines: Indicate equilibrium between phases (fusion, vaporization, sublimation).

• Triple Point: All 3 phases coexist.

• Critical Point: Beyond this, gas and liquid are indistinguishable (supercritical fluid).

6. Analyzing Phase Diagram

• Determine phase at given T and P.

• Predict phase changes with changes in T or P.

• Understand slope of solid-liquid line (e.g., water has negative slope → ice melts under pressure).

7. Types of Solutions

• Solid, liquid, or gas solute in solid, liquid, or gas solvent.

• Common types:

  • Gas in gas (air)

  • Solid in liquid (saltwater)

  • Gas in liquid (soda)

  • Liquid in liquid (alcohol in water)

  • Solid in solid (alloy)

8. The Solution Process (Factors Affecting Solubility)

• Steps:

  1. Solute particles separate.

  2. Solvent particles separate.

  3. Solute and solvent interact.

• Factors:

  • Temperature: Solids ↑ solubility with ↑ T; gases ↓.

  • Pressure: Affects gas solubility (Henry’s Law).

  • Nature of solute & solvent: “Like dissolves like.”

9. Expression of Concentration

• Qualitative:

  • Dilute vs. concentrated.

• Quantitative:

  • Molarity (M): mol solute / L solution

  • Molality (m): mol solute / kg solvent

  • % by mass / volume

  • Parts per million (ppm) / billion (ppb)

  • Mole fraction (χ): mol component / total mol

10. Preparation of Solution

• Steps:

  1. Calculate required amount of solute.

  2. Dissolve solute in some solvent.

  3. Transfer to volumetric flask.

  4. Add solvent to the mark.

• Use M₁V₁ = M₂V₂ for dilutions.

11. Solution Stoichiometry

• Combines molarity with chemical equations.

• Process:

  1. Write balanced equation.

  2. Use volume and molarity to find moles.

  3. Use mole ratio to find unknown.

  4. Convert moles back to required unit (volume, mass, etc.).

• Definition: Properties that depend on the number of solute particles, not their identity.

• Only applicable to nonvolatile, nonelectrolyte or electrolyte solutes in a solution.

A. Osmotic Pressure (Π)

• Pressure required to stop osmosis (flow of solvent into solution).

• Formula:
  Π = MRT

  • M = molarity

  • R = gas constant (0.0821 L·atm/mol·K)

  • T = temperature in Kelvin

• Higher solute concentration → higher osmotic pressure.

B. Vapor Pressure Lowering

• Adding solute decreases the vapor pressure of the solvent.

• Raoult’s Law:
  P_solution = X_solvent × P_pure solvent

  • X = mole fraction

  • Less surface area for evaporation → lower vapor pressure.

C. Boiling Point Elevation (ΔTᵦ)

• Solution boils at higher temperature than pure solvent.

• Formula:
  ΔTᵦ = iKᵦm

  • i = van’t Hoff factor (number of particles)

  • Kᵦ = boiling point elevation constant

  • m = molality

• New boiling point = T_b(pure) + ΔTᵦ

D. Freezing Point Depression (ΔT𝒻)

• Solution freezes at lower temperature than pure solvent.

• Formula:
  ΔT𝒻 = iK𝒻m

  • K𝒻 = freezing point depression constant

• New freezing point = T_f(pure) − ΔT𝒻

2. Concept of Nuclear Reactions

• Involve changes in the nucleus of an atom.

• Unlike chemical reactions, nuclear reactions can change elements.

A. Types of Nuclear Reactions

• Alpha Decay (α):

  • Emits ⁴₂He nucleus

  • Decreases mass by 4, atomic number by 2

• Beta Decay (β⁻):

  • Neutron → proton + electron

  • Electron (β⁻) is emitted

• Positron Emission (β⁺):

  • Proton → neutron + positron

• Gamma Decay (γ):

  • Releases energy, no change in mass or atomic number

B. Fission

• Heavy nucleus splits into smaller nuclei + energy (e.g., uranium-235).

• Chain reactions possible.

• Used in nuclear reactors and bombs.

C. Fusion

• Two light nuclei combine into a heavier nucleus.

• Releases much more energy than fission.

• Occurs in stars (e.g., hydrogen fusion into helium).

D. Radioactive Decay

• Unstable isotopes release particles to become stable.

• Measured by half-life: time for half the sample to decay.