topic 6-rate of chemical change

• Rate of reaction


  How fast the reactants are changed into products
• Rates of reaction are pretty important in the chemical industry
• The faster you make chemicals, the more money you make
• Rates of reaction


  • Rusting of iron (slow)
  • Chemical weathering like acid rain damage to limestone buildings (moderate speed)
  • Magnesium reacting with acid to produce bubbles (moderate speed)
  • Explosions (fast)
• Reaction rate


  The speed of a chemical reaction, measured by the amount of product formed or reactant used up over time
• Measuring reaction rate


  1. Record amount of product formed

  2. Record amount of reactant used up

  3. Plot graph of amount vs time
• Steepness of graph line

  Faster the rate of reaction
• Reaction progresses


  Graph line becomes less steep as reactants are used up
• Graphs showing reaction rate

  • Graph 1 (original reaction)
  • Graph 2 and 3 (same initial reactants, faster reaction)
  • Graph 4 (more reactants added, faster reaction)
• Collision theory


  • Particles must collide with enough energy to react
  • Increasing collision frequency increases rate
• Activation energy


  Minimum energy particles need to react and start the reaction
• Factors affecting reaction rate

  • Temperature
  • Concentration/pressure
  • Surface area
  • Catalyst
• How factors increase reaction rate


  1. Increase particle collision frequency

  2. Increase particle collision energy
• Catalyst


  Substance that speeds up a reaction without being consumed
• Enzyme


  Biological catalyst that catalyses reactions in living things
• Measuring reaction rate


  1. Precipitation and colour change

  2. Change in mass (gas given off)

  3. Volume of gas given off
• Measuring reaction rate


  • Magnesium and HCl reaction producing H2 gas
• Increasing acid concentration


  Increases reaction rate
• Magnesium and HCl React to Produce H2 Gas


  1. Add dilute hydrochloric acid to conical flask

  2. Add magnesium ribbon

  3. Start stopwatch and record mass at regular intervals

  4. Plot results in table and graph

  5. Repeat with more concentrated acid solutions
• Putting cotton wool in the top of the flask lets the gas escape but stops the acid spitting out
• Variables such as the amount of magnesium ribbon and the volume of acid used should be kept the same, only change the acid's concentration to make the experiment a fair test
• Higher concentration of acid

  Faster rate of reaction
• Sodium Thiosulfate and HCl Produce a Cloudy Precipitate


  1. Add dilute sodium thiosulfate to conical flask

  2. Add dilute HCl and start stopwatch

  3. Time how long it takes for black cross to disappear

  4. Repeat with different concentrations of reactants
• The reaction releases sulfur dioxide, so the experiment should be carried out in a well-ventilated place
• Higher concentration of HCl


  Quicker the reaction and the less time it takes for the mark to disappear
• This reaction doesn't give a set of graphs, only readings of how long it took till the mark disappeared for each concentration
• Reversible Reaction


  The products can react to form the reactants again
• Reversible Reactions Will Reach Equilibrium


  1. Forward reaction slows down as reactant concentrations fall

  2. Backward reaction speeds up as product concentrations rise

  3. Equilibrium is reached when forward and backward reactions are at the same rate

  4. Equilibrium is only reached in a closed system
• Equilibrium


  • Both reactions are still happening, but there's no overall effect
  • Concentrations of reactants and products have reached a balance and won't change
• Position of Equilibrium


  Can be on the right (greater concentration of products) or the left (greater concentration of reactants)
• Conditions that affect equilibrium position


  • Temperature
  • Pressure (for reactions involving gases)
  • Concentration of reactants and products
• Heating ammonium chloride moves the equilibrium to the right (more ammonia and hydrogen chloride), cooling it moves it to the left (more ammonium chloride)
• Endothermic and Exothermic Reversible Reactions

  • If the reaction is endothermic in one direction, it will be exothermic in the other
  • The energy transferred from the surroundings by the endothermic reaction is equal to the energy transferred to the surroundings during the exothermic reaction
• Heating blue hydrated copper(II) sulfate crystals drives off the water, leaving white anhydrous copper(II) sulfate powder. This is endothermic. Adding water to the white powder gets the blue crystals back, which is exothermic.
• Le Chatelier's Principle


  If you change the conditions of a reversible reaction at equilibrium, the system will try to counteract that change
• Decreasing temperature


  Equilibrium moves in the exothermic direction to produce more heat
• Increasing temperature


  Equilibrium moves in the endothermic direction to try and decrease it
• Increasing pressure


  Equilibrium moves in the direction where there are fewer molecules of gas
• Decreasing pressure


  Equilibrium moves in the direction where there are more molecules of gas
• Increasing concentration of reactants

  System tries to decrease it by making more products