chemistry final study guide

Study Guide: Meet the Mole

1. The Mole and Avogadro’s Number

• Mole (mol): Unit for measuring the amount of substance.

• Avogadro’s Number:
  1 mol=6.022×10^23 particles (atoms, molecules, etc.)

2. Conversions

• Moles ↔ Atoms/Molecules:

Moles × 6.022×10^23 = particles

Particles ÷ 6.022×10^23 = moles

• Moles ↔ Grams: Use molar mass from the periodic table (g/mol).

Moles × molar mass=grams

Grams ÷ molar mass=moles

3. Molar Mass

• Mass of 1 mol of a substance in g/mol

• For compounds: Add atomic masses of all atoms

  • Example:

H2O=2(1.008)+16.00=18.02 g/mol

4. Percent Composition

• Formula:

%=(mass of element in 1 mol/molar mass of compound)×100

Example:
CO₂ → 27.29% C, 72.71% O

5. Determining Chemical Formulas

• Empirical Formula: Lowest whole-number ratio of atoms

• Molecular Formula:

Molecular mass÷Empirical formula mass=Factor

Multiply empirical subscripts by the factor.

6. Problem-Solving Strategy: RADAR

• R: Read the question

• A: Analyze what’s given

• D: Decide on a plan

• A: Apply the plan

• R: Reflect on the answer

Study Guide: Writing Chemical Reactions

1. Types of Changes

• Physical Change: No change in substance identity (e.g., melting, boiling).

• Chemical Change: Produces new substances (e.g., burning, rusting).

2. Chemical Reactions

• Definition: Atoms rearrange to form new substances.

• Collision Theory: Atoms must collide with enough energy to react.

• Evidence of Reaction:

  • Temperature or light change

  • Color change

  • Gas or odor produced

  • Formation of a precipitate (solid)

3. Law of Conservation of Mass

• Matter cannot be created or destroyed.

• Total mass and atom count must be the same before and after a reaction.

Mass of Reactants=Mass of Products

4. Writing Chemical Equations

• Word Form Example: Hydrogen gas reacts with oxygen gas to form water.

• Formula Form Example:

2H2(g)+O2(g)→2H2O(l)

Symbols:

• (s) = solid

• (l) = liquid

• (g) = gas

• (aq) = aqueous (dissolved in water)

5. Parts of an Equation

• Reactants: Starting substances (left side)

• Products: Substances formed (right side)

• Coefficients: Indicate number of molecules (can change)

• Subscripts: Show number of atoms in a molecule (cannot change)

6. Balancing Equations

Steps:

1. Write the unbalanced equation.

2. List and count atoms of each element on both sides.

3. Add coefficients to balance atoms.

4. Keep subscripts unchanged.

5. Use lowest whole number ratio.

Example:

Fe+Cl2→FeCl3 (Balanced: 2Fe+3Cl2→2FeCl3)

Study Guide: Molecular Geometry

1. Molecular Structure Basics

• Chemical Formula – Tells what atoms are in a molecule.
  Example: C₆H₁₂O₆

• Structural Formula – Shows a 2D layout of the atoms.

• Actual molecules are 3D!

2. VSEPR Theory

• VSEPR = Valence Shell Electron Pair Repulsion

• Electron pairs repel each other and spread out as far as possible.

• Electron Domain = Area of electron density (bond or lone pair)

3. Types of Electron Domains

• Bonding pair – Shared between atoms

• Lone pair – Non-bonded electrons

• Multiple bonds (double/triple) count as one domain

4. Molecular Shapes (Based on Electron Domains)

5. 3D Notation

• Molecules are drawn using wedge-and-dash notation to show 3D shape.

  • Solid line = in plane

  • Wedge = out of plane

  • Dash = behind plane

6. Practice Examples

• CH₂O → 3 domains, 0 lone pairs → Trigonal Planar

• CO₂ → 2 domains, 0 lone pairs → Linear

• SCl₂ → 4 domains, 2 lone pairs → Bent

Study Guide: Intermolecular Forces (IMFs)

1. Key Vocabulary

• Inter = between molecules (IMFs)

• Intra = within a molecule (chemical bonds)

2. Polarity

• Polar Molecules: Unequal sharing of electrons → partial charges (e.g. H₂O)

• Nonpolar Molecules: Equal sharing of electrons → no partial charges (e.g. O₂)

• Use:

  • BEND → Bond Electronegativity Difference

    • ΔEN = 0 → nonpolar

    • ΔEN > 0 → polar

  • SNAP:

    • Symmetrical = Nonpolar

    • Asymmetrical = Polar

3. Types of Intermolecular Forces

4. IMFs and Properties

• Boiling & Melting Point ↑ with stronger IMF

• Viscosity ↑ with stronger IMF

• Polar molecules generally have stronger IMFs than nonpolar ones

5. Practice Strategy

For each substance:

• Identify if it’s polar or nonpolar

• Determine the dominant IMF:

  • Nonpolar → London Dispersion

  • Polar with H–N/O/F → Hydrogen Bond

  • Polar without H–N/O/F → Dipole-Dipole

🧪 Chemical Reaction Types – Study Guide

Chemical reactions can be classified into five major types to help predict products:

1. Synthesis (Combination) Reactions

• Definition: Two or more reactants form one product.

• General Formula: A + B → AB

• Examples:

  • 2Na + Cl₂ → 2NaCl

  • C + O₂ → CO₂

  • CaO + SO₂ → CaSO₃ (a salt)

• Also called: Composition, Combination, or Formation Reactions

2. Decomposition Reactions

• Definition: A single compound breaks into two or more simpler substances.

• General Formula: AB → A + B

• Examples:

  • Ca(OH)₂ → CaO + H₂O

  • 2H₂O → 2H₂ + O₂ (electrolysis with electricity)

• Requires: Energy (heat or electricity)

3. Combustion Reactions

• Definition: A substance reacts with oxygen, often producing heat/light.

• General Formula: Reactant + O₂ → Products

• Examples:

  • 2H₂ + O₂ → 2H₂O

  • CH₄ + 2O₂ → CO₂ + 2H₂O (hydrocarbon combustion)

4. Single Replacement (Displacement) Reactions

• Definition: One element replaces a similar element in a compound.

• General Formula: A + BC → B + AC

• Examples:

  • 2Al + 3Pb(NO₃)₂ → 3Pb + 2Al(NO₃)₃

  • Cl₂ + 2KBr → Br₂ + 2KCl

  • 2Na + 2H₂O → 2NaOH + H₂

• Occurs in: Aqueous solutions

• Key concept: Use the activity series to predict if replacement will occur.

5. Double Replacement (Displacement) Reactions

• Definition: Ions from two compounds switch places.

• General Formula: AB + CD → AD + CB

• Examples:

  • NaCl + AgNO₃ → AgCl (precipitate) + NaNO₃

  • FeS + 2HCl → H₂S (gas) + FeCl₂

  • HCl + NaOH → NaCl + H₂O (molecular compound)

⚡ Predicting Products – Activity Series

• A list showing reactivity of elements.

• More reactive metals: lose electrons more easily.

• More reactive nonmetals: gain electrons more easily.

• Rules: An element can replace another below it in the series, but not one above.

🧲 Ionic Bonds – Study Guide

🔹 What are Ionic Compounds?

• Made of positive (cations) and negative (anions) ions.

• Formed when electrons are transferred from a metal (low electronegativity) to a nonmetal (high electronegativity).

• The goal: overall charge = 0 (neutral compound).

🔹 Ions

• Cations: Positive ions (lose electrons); usually metals.
  Ex: Li → Li⁺

• Anions: Negative ions (gain electrons); usually nonmetals.
  Ex: N → N³⁻

• Oxidation Numbers: Show the ion's charge; follow group trends on the periodic table.

🔹 Rule of Zero Charge

• The total positive and negative charges in a compound must cancel out.

🔹 Using Lewis Structures

1. Draw individual element diagrams.

2. Show electrons being transferred (arrows).

3. Add atoms until all are stable.

4. Write the chemical formula (using subscripts).

5. Double-check: total charge = 0.

🔹 Writing Formulas (Name → Formula)

1. Write ion symbols with charges.

2. Crisscross charges to subscripts.

3. Write formula; simplify if needed.

Example:
Calcium Nitride
Ca²⁺, N³⁻ → Ca₃N₂

🔹 Naming Compounds (Formula → Name)

1. Name the metal first.

2. Name the nonmetal with “-ide” ending.
   Ex: NaCl → Sodium chloride

🔹 Exceptions

1. Polyatomic Ions

• Groups of atoms acting as a single ion (e.g., SO₄²⁻, NH₄⁺).

• Don’t change their subscripts.

• Use special names (e.g., NaNO₃ → Sodium nitrate).

2. Transition Metals

• Can have multiple charges (e.g., Fe²⁺, Fe³⁺).

• Use Roman numerals to show charge:
  FeCl₃ → Iron (III) chloride

✅ Tips

• Use the Periodic Table to predict charges.

• Always check for neutrality.

• For polyatomic ions: use parentheses and don't alter the group.

⚖ Chemical Equilibrium – Study Guide

🔁 What Is Chemical Equilibrium?

• Occurs in reversible reactions:
  Example:

  • Forward: 2SO₂ + O₂ → 2SO₃

  • Reverse: 2SO₃ → 2SO₂ + O₂

• Equilibrium is when the forward and reverse reactions happen at the same rate.

• Represented by a double arrow (⇌).

• The concentrations remain constant, not necessarily equal.

📐 Le Chatelier’s Principle

If a system at equilibrium is disturbed, it will shift to counteract the disturbance and restore equilibrium.

🔄 Factors That Affect Equilibrium

1. Concentration

• Add Reactant → Shift toward products

• Remove Reactant → Shift toward reactants

• Add Product → Shift toward reactants

• Remove Product → Shift toward products

2. Pressure (gases only)

• Increase Pressure → Shift to the side with fewer gas molecules

• Decrease Pressure → Shift to the side with more gas molecules

3. Real-Life Example (Breathing & CO₂ levels)

• Exercise adds CO₂ → Shift to reactants

• Rapid breathing removes CO₂ → Shift to products