Chapter 5

• Exothermic reactions are reactions that release energy in the form of heat or light.
• Endothermic reactions absorb energy from their surroundings.
• Exothermic example:

  • octane + oxygen gas → carbon dioxide + water vapour + energy
• Reactions with oxygen are exothermic.
• Combustion reactions are any chemical reaction in which a substance burns in oxygen gas to produce light and heat. e.g. lighting a match or burning gas on a stove.
• Combustion Example

  • Hydrogen gas + oxygen gas → water vapour
  • Magnesium + oxygen → magnesium oxide
• Corrosion/rusting occurs when metals combine with oxygen gas to form a metal oxide.


  Metal + oxygen → metal oxide
• Corrosion example

  • Rusting of iron: iron + oxygen + water → hydrated iron (iii) oxide
• Respiration is the combustion of glucose with oxygen to produce carbon dioxide and water.
• Activity Series shows how easily metals react with oxygen, water and acid.
• How fast a metal reacts is called reactivity.
• Activity Series

  • Potassium Sodium Calcium Magnesium Aluminum Zinc Iron Nickel Tin Lead Copper Silver Gold.
• Chemical ice packs use an endothermic reaction between ammonium chloride and water to make the pack cold.
• Decomposition reaction is when a reactant breaks apart to form several products, the reactant decomposes.


  XY → X + Y
• Decomposition Example

  Carbonic acid → water + carbon dioxide
• Thermal decomposition is when substances decompose when heated.
• Metal carbonates and metal hydrogen carbonates undergo thermal decomposition when heated.
• Thermal Decomposition Example

  1. Sodium hydrogen carbonate → sodium carbonate + carbon dioxide + water
  2. Airbags: sodium azide → sodium + nitrogen gas
• Combination reactions occur when two reactants combine to form a single product.

  • X + Y → XY
• Combination Example

  • Hydrogen + Chloride → Hydrogen Chloride
• Precipitation reactions occur when two soluble reactants combine to form an insoluble product called a precipitate.
• When a soluble substance is dissolved, the particles are spread thinly throughout the solution.
• Ionic compounds are substances made up of a crystal lattice of positive ions (cations) and negative ions (anions).
• In an ionic compound the cation comes first then the anion.
• Ionic compounds are neutral, total charge on cations balances the total charge on anions.
• Acids

  • Substances that release H+ ions.
  • pH of less than 7
  • Corrosive
• Examples of Acids

  • Hydrochloric acid (HCl)
  • Sulfuric Acid
  • Nitric Acid
• Bases

  • Release OH- ions (hydroxide ions)
  • pH of more than 7
  • Alkaline solutions
  • Caustic
• Examples of Bases

  • Sodium Hydroxide (NaOH)
  • Calcium Hydroxide
  • Sodium Bicarbonate (baking soda)
• Types of acid reactions

  • Neutralisation
  • Acid - Metal
  • Acid - Carbonate
• Neutralisation is when an acid reacts with a base.
• In neutralisation, the H+ ions from acid bonds with OH- ions from base to form water.
• Neutralisation

  • Acid + Base → Salt + water
• Neutralisation

  • Sulfuric acid + magnesium hydroxide → Magnesium sulfate + water
• Everyday Neutralisation

  • Heartburn and indigestion are caused by excess acid.
  • Antacids contain the base Magnesium Hydroxide to neutralise the acid.
  • Hydrochloric acid + Magnesium hydroxide → Magnesium chloride + water
• Acid - Metal

  • acid + metal → salt + hydrogen gas
• Acid Metal Reactions

  • hydrochloric acid + magnesium → magnesium chloride + hydrogen gas
  • Sulfuric acid + iron → iron (iii) sulfate + hydrogen gas
• Acid - carbonate

  • Acid + Carbonate → salt + water + carbon dioxide
• Acid - Carbonate Reactions

  • Hydrochloric acid + calcium carbonate → calcium chloride + water + carbon dioxide
  • Sulfuric acid + calcium carbonate → calcium sulfate + water + carbon dioxide
• Rate of reaction is the speed at which a chemical reaction proceeds.