Types of Chemical Reactions

Introduction

• Overview of different reactions in chemistry class:
  • Combination (Synthesis) Reactions
  • Decomposition Reactions
  • Combustion Reactions
  • Single Replacement Reactions
  • Double Replacement Reactions
  • Precipitation Reactions
  • Acid-Base Reactions
  • Gas Evolution Reactions
  • Redox (Redux) Reactions

Combination (Synthesis) Reactions

• Definition: Two smaller substances combine to form a larger substance.
• Example: Magnesium metal (Mg) reacts with Nitrogen gas (N2) to form Magnesium Nitride (Mg3N2).
  • Use crisscross method to determine formula: Mg^(2+) and N^(3-) → Mg3N2.
  • Balance the reaction using coefficients.
  • This reaction is also a Redox reaction.
• Redox Details:
  • Oxidation: Magnesium (0 to +2)
  • Reduction: Nitrogen (0 to -3)
  • Oxidizing Agent: Substance reduced (Nitrogen)
  • Reducing Agent: Substance oxidized (Magnesium)
• General Tip: If a pure element is on one side of the reaction and in a compound on the other, it is likely a Redox reaction.

Additional Examples of Combination Reactions

• Calcium and Oxygen: Forms Calcium Oxide (CaO)
  • Use of crisscross method.
  • Balanced by: 2 Ca + O2 → 2 CaO.
• Aluminum and Bromine: Forms Aluminum Bromide (AlBr3)
  • Balancing with least common multiple method.
  • General strategy for balancing such reactions.

Decomposition Reactions

• Definition: A compound (AB) breaks down into simpler components (A + B).
• Examples:
  • Carbonic Acid (H2CO3) decomposes to Water (H2O) and CO2.
  • Sulfuric Acid (H2SO4) heated to produce Sulfur Trioxide (SO3) and Water.
  • Electrolysis of Water into Hydrogen (H2) and Oxygen (O2).
• Redox Example: Water decomposition
  • Hydrogen reduced (H (+1) to H2 (0))
  • Oxygen oxidized (O (-2) to O2 (0))
  • Requires energy input (non-spontaneous)
• Common Decomposition Reactions:
  • Heating Mercury Oxide (HgO): Produces Mercury and Oxygen.
  • Heating metal carbonates: Produces CO2 and metal oxides.
  • Heating metal hydroxides: Produces water and metal oxides.

Combustion Reactions

• Definition: Hydrocarbon reacts with oxygen producing CO2 and water, releasing energy.
• Example: Methane (CH4) + O2 → CO2 + H2O.
  • Always a Redox reaction.
• Balancing Method: Balance carbon, then hydrogen, then oxygen.
• Additional Example: Ethane (C2H6) combustion
  • Follow similar balancing steps.

Single Replacement Reactions

• Definition: A pure element displaces another element in a compound.
• Example: Aluminum + Copper Chloride → Aluminum Chloride + Copper.
  • Use activity series to determine if reaction occurs.
  • Aluminum more active than Copper.
• Balancing and Phases: Adjust coefficients and assign phases.
• Redox Details:
  • Oxidation: Aluminum (0 to +3)
  • Reduction: Copper (+2 to 0)

Double Replacement Reactions

• Definition: Two compounds exchange ions to form two new compounds.
• Example: Silver Nitrate + Magnesium Chloride → Silver Chloride + Magnesium Nitrate
  • Solubility rules to determine phases.
  • Precipitation Reaction: Formation of insoluble product (e.g., AgCl).
• Non-Redox Reaction: No pure elements involved.
• Net Ionic Equations: Focus on reactive species only.
  • Example: Ag⁺ + Cl⁻ → AgCl (s)

Acid-Base Reactions

• Definition: Acid and base react to form salt and water.
• Example: Sulfuric acid (H2SO4) and Sodium Hydroxide (NaOH).
  • Balancing: H+ + OH- → H2O.
• Net Ionic Equation: Primary ions involved in forming water.

Gas Evolution Reactions

• Definition: Reaction produces a gas as one of the products.
• Example: Ammonium Chloride + Potassium Hydroxide → Ammonia gas + Water + Potassium Chloride
  • Recognize intermediate formation and decomposition.
• Other Examples:
  • Sodium Carbonate reacts with HCl to produce CO2 gas.

Conclusion

• Reviewed common reaction types with emphasis on Redox distinctions and balancing techniques.
• Understanding these reactions is crucial for mastering chemistry.