chem

Chemistry 3 — January Exam Study Guide (Ultra‑Detailed with Examples)

This guide includes every topic from your exam review, plus important details, worked examples, common mistakes, and memory tips. Use this as your primary study source.

CHAPTER 2 — MATTER

1. States of Matter (What You MUST Know)

Solid

• Definite shape and definite volume

• Particles vibrate in fixed positions

• Strong intermolecular forces

• Example: ice, metal, wood

Liquid

• Definite volume, no definite shape

• Particles are close but can slide

• Medium intermolecular forces

• Example: water, oil

Gas

• No definite shape or volume

• Particles move freely and spread out

• Weak intermolecular forces

• Example: air, oxygen

Exam Tip: If particles are far apart → gas. If sliding → liquid. If vibrating → solid.

2. Physical vs. Chemical Changes (WITH EXAMPLES)

Physical Change

• Does NOT change the identity of the substance

• Only affects form or state

• Usually reversible

Examples:

• Melting copper → still copper

• Ripping paper → still paper

• Boiling water → still H₂O

Chemical Change

• Produces a NEW substance

• Atoms rearrange

• Usually irreversible

Examples:

• Wood burning → ash + gas

• Rusting iron → iron oxide

• Hydrogen + oxygen → water

Exam Tip: If bonds break/form → chemical change.

3. Pure Substances vs. Mixtures

Pure Substance

• Fixed composition

• One type of particle

• Cannot be separated physically

Types:

• Element (Fe, O₂)

• Compound (H₂O, NaCl)

Mixture

• Variable composition

• Two or more substances mixed physically

• Can be separated physically

Types:

• Homogeneous: uniform (air, coffee)

• Heterogeneous: non‑uniform (salad dressing)

4. Classification Examples (EXAM STYLE)

• Chocolate chip cookies → heterogeneous mixture

• Air → homogeneous mixture

• Bronze → homogeneous mixture (alloy)

• Iron metal → element

• Sodium chloride → compound

• Oil & vinegar → heterogeneous mixture

• Coffee → homogeneous mixture

5. Separation Techniques

Filtration

• Separates solid from liquid

• Example: sand + water

Distillation

• Separates liquids by boiling point

• Example: alcohol + water

CHAPTER 3 — MEASUREMENTS & CALCULATIONS

1. Scientific Notation (STEP‑BY‑STEP)

Rules:

• One nonzero digit before decimal

• Power of 10 shows decimal movement

Examples:

• 23400 → 2.34 × 10⁴

• 0.000168 → 1.68 × 10⁻⁴

2. Significant Figures (VERY IMPORTANT)

Rules:

1. Nonzero digits count

2. Zeros between numbers count

3. Leading zeros do NOT count

4. Trailing zeros count only if decimal is present

Examples:

• 0.02398 → 4 sig figs

• 0.0103 → 3 sig figs

3. Sig Figs in Calculations

Multiply / Divide → least sig figs

Example:
2.5 × 3.42 = 8.6 (2 sig figs)

Add / Subtract → least decimal places

4. Conversions (DIMENSIONAL ANALYSIS)

Always:

• Write given

• Multiply by conversion factor

• Cancel units

Temperature:

• °C → K = +273

• K → °F = (K − 273) × 9/5 + 32

5. Density

Formula:
Density = mass ÷ volume

Example:
23.2 g ÷ 25.5 mL = 0.910 g/mL

CHAPTER 4 — CHEMICAL FOUNDATIONS

1. Atomic Scientists (MEMORIZE)

• Electron → J.J. Thomson

• Plum Pudding Model → Thomson

• Proton → Rutherford

• Nucleus → Rutherford

• Gold Foil Experiment → Rutherford

• Neutron → James Chadwick

2. Isotopes

• Same element

• Same protons

• Different neutrons

Example:
Carbon‑12 vs Carbon‑14

3. Counting Subatomic Particles

Rules:

• Protons = atomic number

• Electrons = protons (neutral)

• Neutrons = mass − atomic number

4. Periodic Table Groups

• Group 1: Alkali Metals (1+ charge)

• Group 2: Alkaline Earth Metals (2+)

• Groups 3–12: Transition Metals

• Group 17: Halogens (−1)

• Group 18: Noble Gases (no charge)

5. Diatomic Molecules

HOFBrINCl
H₂, O₂, F₂, Br₂, I₂, N₂, Cl₂

6. Ions

• Ion: charged atom

• Cation: positive (lost electrons)

• Anion: negative (gained electrons)

Metals LOSE electrons → cations
Nonmetals GAIN electrons → anions

CHAPTER 5 — NOMENCLATURE

1. Compound Types

• Type I: metal + nonmetal (fixed charge)

• Type II: metal + nonmetal (variable charge)

• Type III: nonmetal + nonmetal

2. Naming Rules (WITH EXAMPLES)

Type I
NaCl → sodium chloride

Type II
FeS → iron (II) sulfide

Type III
CO₂ → carbon dioxide
N₂O₅ → dinitrogen pentoxide

3. Acids

Binary:
HCl → hydrochloric acid

Oxyacids:
-ate → ic (nitrate → nitric)
-ite → ous (nitrite → nitrous)

CHAPTER 6 — CHEMICAL COMPOSITION

1. Avogadro’s Number

6.022 × 10²³ particles/mol

2. Molar Mass

Add atomic masses
Units: g/mol

Example:
CaCO₃ = 100.09 g/mol

3. Mole Conversions (ALL TYPES)

Use factor‑label method

g ↔ mol ↔ particles

4. Percent Composition

Formula:
(element mass ÷ total mass) × 100

5. Empirical vs Molecular

Empirical = simplest ratio
Molecular = actual formula

Steps:

1. Molar mass ÷ empirical mass

2. Multiply subscripts

CHAPTER 7 — BALANCING REACTIONS

1. Physical States

(s) solid
(l) liquid
(g) gas
(aq) aqueous

2. Conservation of Mass

Atoms are never created or destroyed

3. Balancing Rules

• Add coefficients only

• Never change subscripts

• Balance metals → nonmetals → H → O last

4. Writing Equations

Steps:

1. Write correct formulas

2. Balance atoms

3. Add states

FINAL EXAM STRATEGY

• Always show work

• Watch sig figs

• Units must cancel

• Double‑check charges

• Recount atoms after balancing

This guide now includes definitions, rules, examples, and exam tips for every topic on your review sheet.

CHAPTER 2 — MATERIA AND CHANGE

1. States of Matter & Phase Changes

Solid

• Definite shape and volume.

• Particles vibrate in fixed positions due to high density and maximum intermolecular forces.

• Phase Changes: Melting (Solid (\rightarrow) Liquid), Sublimation (Solid (\rightarrow) Gas).

Liquid

• Definite volume, no definite shape (takes shape of container).

• Particles are close but fluid; they slide past one another.

• Phase Changes: Freezing (Liquid (\rightarrow) Solid), Evaporation/Boiling (Liquid (\rightarrow) Gas).

Gas

• No definite shape or volume; highly compressible.

• Particles move randomly at high speeds with minimal intermolecular attraction.

• Phase Changes: Condensation (Gas (\rightarrow) Liquid), Deposition (Gas (\rightarrow) Solid).

2. Physical vs. Chemical Properties

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

  • Intensive: Independent of amount (e.g., density).

  • Extensive: Dependent on amount (e.g., mass, volume).

• Chemical Properties: Ability to undergo changes that transform it into different substances (e.g., flammability, reactivity with acid).

3. Classification of Matter

• Pure Substances: Constant composition.

  • Elements: Simplest form of matter (e.g., $Au$, $H_{2}$).

  • Compounds: Chemically bonded elements in fixed ratios (e.g., $NaCl$, $H_{2}O$).

• Mixtures: Physical blend of two or more substances.

  • Homogeneous (Solutions): Uniform throughout (e.g., salt water, alloys like brass).

  • Heterogeneous: Distinct phases visible (e.g., sand in water, granite).

CHAPTER 3 — MEASUREMENTS & CALCULATIONS

1. Accuracy vs. Precision

• Accuracy: How close a measurement is to the true/accepted value.

• Precision: How close a series of measurements are to one another (reproducibility).

2. Significant Figures (Sig Figs)

• Rule 1: Non-zero digits are always significant.

• Rule 2: Captive zeros (between non-zeros) are significant (e.g., $405$ has 3).

• Rule 3: Leading zeros are NEVER significant (e.g., $0.002$ has 1).

• Rule 4: Trailing zeros are significant ONLY if there is a decimal point (e.g., $100.$ has 3, $100$ has 1).

3. Calculations with Sig Figs

• Multiplication/Division: Round to the fewest total sig figs.

• Addition/Subtraction: Round to the fewest decimal places.

4. Density & Energy

• Density: $\text{Density} = \frac{\text{mass}}{\text{volume}}$. Floating occurs if \text{D}{object} < \text{D}{liquid}.

• Specific Heat Capacity: $q = m \cdot c \cdot \Delta T$

  • $q$ = heat (Joules)

  • $m$ = mass (grams)

  • $c$ = specific heat

  • $\Delta T$ = change in temperature ($T<em>{final} - T</em>{initial}$)

CHAPTER 4 — ATOMIC STRUCTURE

1. Atomic Theory Timeline

• Dalton: Atomic Theory (atoms are indivisible spheres).

• Thomson: Discovered electrons using Cathode Ray Tube; proposed "Plum Pudding" model.

• Rutherford: Gold Foil Experiment. Discovered the nucleus (dense, positive center) and that the atom is mostly empty space.

• Bohr: Electrons exist in specific energy levels/orbits.

• Chadwick: Discovered the neutron.

2. Isotopes & Atomic Mass

• Isotopes: Atoms of the same element with the same number of protons but different numbers of neutrons. This affects the mass number.

• Average Atomic Mass: The weighted average of all naturally occurring isotopes.

  • $\text{Avg Mass} = (\text{mass}<em>{1} \times \%</em>{1}) + (\text{mass}<em>{2} \times \%</em>{2}) \dots$

3. Periodic Table Organization

• Periods: Horizontal rows (represent energy levels).

• Groups/Families: Vertical columns (elements share similar chemical properties due to valence electrons).

  • Group 1: Alkali Metals (highly reactive, $1+$ charge).

  • Group 2: Alkaline Earth Metals ($2+$ charge).

  • Group 17: Halogens (highly reactive non-metals, $1-$ charge).

  • Group 18: Noble Gases (inert, stable octet).

CHAPTER 5 — NOMENCLATURE & POLYATOMIC IONS

1. Common Polyatomic Ions (Memorize!)

• Nitrate: $NO_{3}^{-1}$

• Sulfate: $SO_{4}^{-2}$

• Carbonate: $CO_{3}^{-2}$

• Phosphate: $PO_{4}^{-3}$

• Ammonium: $NH_{4}^{+1}$

• Hydroxide: $OH^{-1}$

2. Naming Systems

• Binary Ionic: Name metal + nonmetal-ide (e.g., $MgCl_{2}$ = magnesium chloride).

• Transition Metals (Type II): Use Roman Numerals for the charge (e.g., $FeCl_{3}$ = iron (III) chloride).

• Molecular (Type III): Use prefixes (mono-, di-, tri-, tetra-, etc.).

  • Example: $P<em>{2}O</em>{5}$ = diphosphorus pentoxide.

CHAPTER 6 — CHEMICAL COMPOSITION

1. The Mole Concept

• 1 mole = $6.022 \times 10^{23}$ particles (Avogadro's Number).

• Molar Mass: Sum of atomic masses from the periodic table in g/mol.

2. Empirical vs. Molecular Formulas

• Empirical: Lowest whole-number ratio (e.g., $CH_{2}O$).

• Molecular: The actual formula (e.g., $C<em>{6}H</em>{12}O_{6}$).

• Calculation Steps for Empirical:

  1. Percent to mass (assume 100g).

  2. Mass to moles (divide by molar mass).

  3. Divide by small (divide all mole values by the smallest number).

  4. Multiply 'til whole (if necessary).

CHAPTER 7 — CHEMICAL REACTIONS

1. Evidence of a Reaction

• Color change, evolution of a gas (bubbles), formation of a precipitate (solid), or temperature change (release/absorption of energy).

2. Classifying Reactions

• Synthesis: $A + B \rightarrow AB$

• Decomposition: $AB \rightarrow A + B$

• Single Replacement: $A + BC \rightarrow AC + B$

• Double Replacement: $AB + CD \rightarrow AD + CB$

• Combustion: $\text{Hydrocarbon} + O<em>{2} \rightarrow CO</em>{2} + H_{2}O$

3. Balancing Equations

• Use Coefficients to ensure the number of atoms on the reactant side equals the product side (Law of Conservation of Mass).