Chemistry Exam Summary Notes

Reactivity Series

  • The reactivity series ranks metals based on their ability to lose electrons and react with acids, oxygen, and water.
  • Order of Reactivity (Most to Least):
    1. Highly reactive metals: Potassium (K), Sodium (Na), Calcium (Ca)
      • React violently with water and acids.
    2. Moderately reactive metals: Magnesium (Mg), Zinc (Zn), Iron (Fe)
      • React readily with acids, but slower with water.
    3. Least reactive metals: Copper (Cu), Silver (Ag), Gold (Au)
      • Do not react with acids or water.

Displacement Reactions

  • A more reactive metal can replace a less reactive metal from its compound.
  • Example: Zinc placed in copper sulfate solution → Zinc sulfate + Copper.
    • Zn + CuSO₄ → ZnSO₄ + Cu

Rate of Reaction

  • The rate of reaction tells us how fast a chemical reaction occurs.

Factors Affecting Reaction Rate

  1. Temperature:
    • Higher temperature increases kinetic energy, causing particles to collide more frequently.
  2. Concentration:
    • More reactant particles = higher chance of successful collisions.
  3. Surface Area:
    • More exposed particles lead to faster reactions (e.g., powdered reactants react faster than lumps).
  4. Catalysts:
    • Speed up reactions without being consumed.
  5. Pressure (for gases):
    • Higher pressure forces particles closer together, increasing reaction rate.
  • Example: Magnesium reacting with hydrochloric acid is faster at higher temperatures because molecules move faster.

Extraction of Metals

  • Metals are extracted from ores based on their reactivity.

Methods of Extraction

  1. Highly reactive metals (e.g., Al, Na, K) → Electrolysis.
  2. Moderately reactive metals (e.g., Zn, Fe, Pb) → Reduction using carbon.
  3. Low-reactivity metals (e.g., Cu, Ag, Au) → Found as free elements or extracted using roasting.

Blast Furnace (Iron Extraction)

  • Used to extract iron from its ore (Haematite - Fe₂O₃).
  • Reaction Process:
    • Carbon (C) reacts with oxygen → CO₂
    • CO₂ reacts with more carbon → Carbon monoxide (CO)
    • CO reduces iron ore to pure iron:
      • Fe₂O₃ + 3CO → 2Fe + 3CO₂

Salts and Their Preparation

What Is a Salt?

  • A salt is formed when an acid reacts with a base.
  • It consists of:
    • Cation (Positive ion from the base).
    • Anion (Negative ion from the acid).

Solubility of Salts

Soluble Salts:

  • All sodium, potassium, and ammonium salts.
  • All nitrates.
  • Most chlorides, except silver chloride and lead chloride.
  • Most sulfates, except calcium sulfate, barium sulfate, and lead sulfate.

Insoluble Salts:

  • Silver chloride, lead chloride.
  • Barium sulfate, calcium sulfate, lead sulfate.
  • Most carbonates, except sodium, potassium, and ammonium carbonates.

Methods of Salt Preparation

  1. Neutralization (Acid + Base)
    • Acid + Base → Salt + Water
    • The type of acid determines the type of salt:
      • Sulfuric acid → Sulfate.
      • Hydrochloric acid → Chloride.
      • Nitric acid → Nitrate.
  2. Acid + Metal Reactions
    • Acid + Metal → Salt + Hydrogen
    • Example: Sulfuric acid + Zinc → Zinc sulfate + Hydrogen gas.
  3. Acid + Metal Oxide Reactions
    • Acid + Metal Oxide → Salt + Water
    • Example: Hydrochloric acid + Copper oxide → Copper chloride + Water.
  4. Acid + Carbonate Reactions
    • Acid + Carbonate → Salt + Water + Carbon Dioxide
    • Example: Nitric acid + Calcium carbonate → Calcium nitrate + CO₂ + Water.
  5. Making Salts Using Alkalis (Titration Method)
    • Alkalis are soluble bases containing hydroxide ions (OH⁻ (aq)).
    • Key Alkalis:
      • Strong: Sodium hydroxide (NaOH), Potassium hydroxide (KOH).
      • Weak: Ammonium hydroxide (NH₄OH).

Titration Steps:

  1. Add acid to a conical flask.
  2. Add indicator (e.g., phenolphthalein).
  3. Slowly add alkali from a burette until the indicator changes color.
  4. Repeat without an indicator for pure salt formation.
  5. Evaporate water to allow crystallization.
  • Example Reaction: HCl + NaOH → NaCl + H₂O

Ammonium Salts & Fertilizers

  • Ammonia is soluble in water, forming ammonium hydroxide.
  • Ammonium salts are used in fertilizers.
  • Example:
    • Ammonia + Hydrochloric acid → Ammonium chloride
    • NH₃ + HCl → NH₄Cl

Preparation of Insoluble Salts (Precipitation Method)

  • Insoluble salts are prepared using precipitation.
  • Method: Mix two solutions containing necessary ions → Insoluble salt forms as precipitate.
  • Example: Silver chloride preparation:
    • Silver nitrate + Sodium chloride → Silver chloride (White precipitate).

Summary Table of Salt Preparation Methods

Reaction TypeExample ReactionProduct
Acid + MetalHCl + Zn → ZnCl₂ + H₂Salt + Hydrogen
Acid + OxideH₂SO₄ + CuO → CuSO₄ + H₂OSalt + Water
Acid + CarbonateHNO₃ + CaCO₃ → Ca(NO₃)₂ + CO₂ + H₂OSalt + Water + CO₂
Alkali + AcidNaOH + HCl → NaCl + H₂OSalt + Water

Types of Chemical Reactions

  • Chemical reactions are classified into different types based on how reactants interact to form products.
  • The main types covered are:
    1. Synthesis (Combination) Reactions
    2. Decomposition Reactions
    3. Single Displacement Reactions
    4. Double Displacement Reactions
    5. Combustion Reactions

1. Synthesis (Combination) Reactions

  • Definition:
    • Two or more reactants combine to form a single product.
    • General formula: A + B → AB
  • Example Reaction:
    • 2K(s) + Cl₂(g) → 2KCl(s)(Potassium + Chlorine → Potassium chloride)
  • How to Predict Products:
    • Identify the elements or compounds reacting.
    • Use the Criss-Cross method to balance charges.
    • Form a stable ionic compound.
  • Practice Examples:
    • Mg + N₂ → ?
      • Mg²⁺, N³⁻ → Mg₃N₂
    • 2Na(s) + Cl₂(g) → 2NaCl(s)
    • Mg(s) + F₂(g) → MgF₂(s)
    • 2Al(s) + 3F₂(g) → 2AlF₃(s)

2. Decomposition Reactions

  • Definition:
    • A single compound breaks down into two or more simpler substances.
      • General formula: AB → A + B
  • Example Reaction:
    • 2HgO(s) → 2Hg(l) + O₂(g)(Mercury(II) oxide → Mercury + Oxygen gas)
  • Common Decomposition Patterns:
    • Metal carbonates → Metal oxide + CO₂
    • Example: CaCO₃ → CaO + CO₂
    • Metal hydroxides → Metal oxide + Water
    • Example: Cu(OH)₂ → CuO + H₂O
    • Metal chlorates → Metal chloride + Oxygen
    • Example: 2KClO₃ → 2KCl + 3O₂

3. Single Displacement Reactions

  • Definition:
    • One element replaces another in a compound.
    • General formula: AB + C → AC + B
  • Example Reaction:
    • Zn(s) + CuCl₂(aq) → ZnCl₂(aq) + Cu(s)(Zinc replaces Copper in the solution)
  • Key Concepts:
    • Metals lose electrons (Oxidation) → Become positive ions.
    • Non-metals gain electrons (Reduction) → Become negative ions.
  • Practice Examples:
    1. Zn(s) + 2HCl(aq) → ZnCl₂ + H₂(g)
    2. 2NaCl(s) + F₂(g) → 2NaF(s) + Cl₂(g)
    3. 2Al(s) + 3Cu(NO₃)₂(aq) → 3Cu(s) + 2Al(NO₃)₃(aq)

4. Double Displacement Reactions

  • Definition:
    • Ions in two compounds swap places.
    • General formula: AB + CD → AD + CB
  • Example Reaction:
    • AgNO₃(aq) + NaCl(s) → AgCl(s) + NaNO₃(aq)(Silver nitrate + Sodium chloride → Silver chloride + Sodium nitrate)
  • How to Predict Products:
    • First and outer ions combine.
    • Inside ions combine.
  • Practice Examples:
    1. HCl(aq) + AgNO₃(aq) → ?
    2. CaCl₂(aq) + Na₃PO₄(aq) → ?
    3. Pb(NO₃)₂(aq) + BaCl₂(aq) → ?
    4. FeCl₃(aq) + NaOH(aq) → ?
    5. H₂SO₄(aq) + NaOH(aq) → ?

5. Combustion Reactions

  • Definition:
    • A hydrocarbon reacts with oxygen gas, producing carbon dioxide and water.
    • General formula: CxHy + O₂ → CO₂ + H₂O
  • Example Reaction:
    • C₆H₁₂ + O₂ → CO₂ + H₂O(Combustion of hexane)
  • Fire Triangle (Requirements for Combustion):
    1. Fuel (Hydrocarbon)
    2. Oxygen
    3. Ignition Source (Heat/Spark)
  • Key Features:
    • Produces large amounts of heat and energy.
    • Used for burning fuels like octane (C₈H₁₈) in gasoline.
  • Practice Examples:
    1. BaCl₂ + H₂SO₄ → ?
    2. C₆H₁₂ + O₂ → ?
    3. Zn + CuSO₄ → ?
    4. Cs + Br₂ → ?
    5. FeCO3 → ?