Chemical Reactivity

• Collision theory


  Particles must collide with sufficient (kinetic) energy and at the correct orientation / angle for collisions to be successful and therefore result in a reaction
• Factors that affect reaction rate

  • Surface area - increasing the surface area of (solid) reactants increases the rate of reaction, with a simple explanation
  • Concentration - increasing the concentration of (aqueous / solution) reactants increases the rate of reaction, with a simple explanation
  • Temperature - increasing the temperature of the reaction increases the rate of reaction, with a simple explanation
• Acid


  An acidic solution produces H+ ions. Acids taste sour and turn litmus red.
• Base / Alkali


  A basic / alkaline solution produces OH- ions. Bases taste bitter and turn litmus blue.
• Common laboratory acids and bases

  • Acids: HCl, HNO3, H2SO4
  • Bases: NaOH, KOH, Na2CO3
• pH


  Measures how acidic or basic a substance is
• pH ranges

  • pH less than 7 is acidic. The amount / concentration of H+ > OH-
  • pH greater than 7 is basic. The amount / concentration of H+ < OH-
  • pH equal to 7 is neutral. The amount / concentration of H+ = OH-
• Universal indicator


  Links the colour of UI to the pH of a solution and to relative amount / concentration of H+ or OH- ions present
• Neutralisation


  In a neutralisation reaction between an acid and a base, a salt is produced
• Salt


  An ionic compound made up of cations (+) and anions (-) combined in a specific ratio
• Identifying and writing salt formulas

  1. Name a salt given its formula

  2. Write the formulae for a salt given its name / the cation and anion in the salt

  3. Identify and write the formula (including correct charges) of the cation and anion present when given the formula / name of the salt
• Specific Learning Outcome

  Describe and explain how the rate of a reaction is affected by different factors - surface area, concentration, temperature and the use of a catalyst
• Rates of Reaction - theory

  1. Collision theory

  2. Activation energy

  3. Frequency

  4. Rate of reaction
• Collision theory


  • Reactions occur only when collisions between particles are successful / effective
  • Particles must collide with enough kinetic energy to overcome the activation energy
  • Particles must collide with correct orientation / angle
• Activation energy


  Minimum energy required to make a reaction happen

  
• Rate


  Speed, a measure of how quickly a reaction proceeds
• Frequency


  Number of successful collisions per second
• Measuring the rate of a reaction


  1. Measuring how fast the reactants are used up

  2. Measuring how fast products form
• Surface area


  • How many particles are exposed to collisions
  • Increasing surface area increases the rate of reaction
  • Explanation: More particles exposed to collisions, more collisions per second, more successful collisions
• Concentration


  • How many particles per unit volume are available for collisions
  • Increasing concentration increases the rate of reaction
  • Explanation: More particles per unit volume, more collisions per second, more successful collisions
• Temperature


  • Affects the kinetic energy of particles and how fast they move
  • Increasing temperature increases the rate of reaction
  • Explanation: Particles gain kinetic energy and move faster, more collisions per second, more successful collisions due to higher proportion of particles with kinetic energy greater than activation energy

  
• Catalyst


  • Lowers the activation energy required for a reaction to occur
  • Adding a catalyst increases the rate of a reaction
  • Explanation: Lower activation energy means higher proportion of particles have kinetic energy greater than activation energy, more successful collisions

  
• Reaction rate graph


  The rate of the reaction is represented by the gradient of the line<|>The steeper the gradient the faster the rate of reaction
• As the reaction proceeds

  The rate of reaction decreases
• When two reactions are compared on the same graph

  They both finish at the same horizontal line
• Explanation for why the rate of reaction at the start of the two Oamaru stone reactions is different
• Rates of Reaction

  Factors affecting the rate of a reaction: surface area, concentration, temperature, catalyst

  
• Rates of Reaction - examples

  • Example 1
  • Example 2
  • Example 3
  • Example 4
  • Example 5
• Example 1

  1. 3 experiments carried out at room temperature

  2. 5 g calcium carbonate and 100 mL hydrochloric acid used

  3. Balloon on top to collect carbon dioxide gas
• Flask A

  • 2 mol L-1 hydrochloric acid, lumps of calcium carbonate
• Flask B

  • 2 mol L-1 hydrochloric acid, powdered calcium carbonate
• Flask C

  • 4 mol L-1 hydrochloric acid, powdered calcium carbonate
• Increasing concentration


  Increases reaction rate
• Increasing surface area


  Increases reaction rate
• Flask C


  Greater reaction rate than Flask B
• Flask B


  Greater reaction rate than Flask A
• Example 2


  2 new flasks, Flask D in warm water bath, Flask E at room temperature
• Increasing temperature

  Increases reaction rate
• Higher temperature in Flask D

  Greater reaction rate than Flask E
• Catalyst


  Increases reaction rate by providing an alternative reaction pathway with lower activation energy