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

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. $FeCO3 → ?$