Chemistry Exam Summary

Reactivity Series

• Ranks metals by their ability to lose electrons and react with acids, oxygen, and water.
• Order of Reactivity (Most to Least):
  • Highly reactive metals: Potassium (K), Sodium (Na), Calcium (Ca)
  • React violently with water and acids.
  • Moderately reactive metals: Magnesium (Mg), Zinc (Zn), Iron (Fe)
  • React readily with acids, but slower with water.
  • 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 yields Zinc sulfate and Copper.
  • $Zn + CuSO₄ → ZnSO₄ + Cu$

Rate of Reaction

• Defines 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:

1. Carbon (C) reacts with oxygen → $CO₂$
2. $CO₂$ reacts with more carbon → Carbon monoxide (CO)
3. 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.
  • 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:

• $HCl + NaOH → NaCl + H₂O$

1. Ammonium Salts & Fertilizers
   • Ammonia is soluble in water, forming ammonium hydroxide.
   • Ammonium salts are used in fertilizers.
   • Example:
     • Ammonia + Hydrochloric acid → Ammonium chloride
     • $NH₃ + HCl → NH₄Cl$
2. Preparation of Insoluble Salts (Precipitation Method)
   • Insoluble salts are prepared using precipitation.
   • Method: Mix two solutions containing necessary ions, and the insoluble salt forms as a 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 based on how reactants interact to form products.
  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:

1. Identify the elements or compounds reacting.
2. Use the Criss-Cross method to balance charges.
3. Form a stable ionic compound.

Practice Examples:

1. $Mg + N₂ → ?$
   • $Mg²⁺$, $N³⁻$ → $Mg₃N₂$
2. $2Na(s) + Cl₂(g) → 2NaCl(s)$
3. $Mg(s) + F₂(g) → MgF₂(s)$
4. $2Al(s) + 3F₂(g) → 2AlF₃(s)$

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:

1. First and outer ions combine.
2. 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. $FeCO3 → ?$