rates of reaction

• Rate of chemical reaction

  The speed with which the reactants get turned into products
• Measuring the rate of reaction


  1. Measure how fast the reactants are being used up

  2. Measure how fast the products are being formed
• Rate of reaction = quantity of reactants used / time taken
• Rate of reaction = quantity of products formed / time taken
• Slow reaction


  • Rusting of iron (years or decades)
• Typical reaction


  • Reaction between magnesium and acid (gentle stream of hydrogen bubbles)
• Fast reaction


  • Explosions like fireworks (fraction of a second)
• The rate of reaction starts off really fast when there are lots of reactants, then slows down as the reaction progresses and the reactants get used up
• As time increases


  Mass of reactants remaining decreases
• As time increases


  Volume of products produced increases
• Units for rate of reaction


  • Grams/second
  • Centimetres cubed/second
  • Moles/minute
  • Decimetres cubed/second
• The rate of reaction can be calculated at a particular time, not just the average rate throughout the entire reaction
• Factors that affect the rate of chemical reactions


  • Temperature
  • Concentration or pressure
  • Surface area
  • Presence or absence of a catalyst
• Collision theory


  In order for particles to react they have to collide with each other with sufficient energy (activation energy)
• Rate of reaction


  Depends on the amount of energy the particles have and the frequency of the collisions
• As temperature increases


  Particles gain more energy, move faster, collide more frequently, and collide with more energy
• As concentration or pressure increases


  More particles per unit of volume, collisions more frequent, rate of reaction increases
• As surface area increases


  More area for collisions to take place, frequency of collisions increases, rate of reaction increases
• Catalyst


  Substance that speeds up a reaction without being used up
• How a catalyst works

  1. Lowers the activation energy by providing an alternative reaction pathway

  2. Increases the proportion of successful collisions
• Transition metals and enzymes are common sources of catalysts
• Measuring the mean rate of a reaction and the rate of reaction at a specific time

  1. Calculate the rate of reaction by dividing either the amount of reactants used or the amount of product formed over the time taken for that change to occur

  2. Plot the volume of hydrogen produced during the reaction on a graph with time on the x-axis and volume on the y-axis

  3. To calculate the mean rate of reaction over a period, find the amount of hydrogen produced in that period and divide by the time

  4. To calculate the rate of reaction at a specific time, draw a tangent to the curve at that time and find the gradient of the tangent
• The rate of reaction is fastest at the beginning and slows down as the reaction progresses
• The graph plateaus once all of the magnesium has been used up
• Tangent


  A straight line that just touches the curve and has the same gradient as the curve at that point
• Calculating the rate of reaction from a graph of magnesium remaining over time


  1. Find the one minute point on the curve

  2. Draw a tangent at that point

  3. Use the change in magnesium and the change in time from the tangent to calculate the rate of reaction
• Examiners will allow answers within a small range of values when judging tangents
• Reversible reaction


  A reaction with a double arrow in the middle, indicating it can react in both forward and backward directions
• Reversible reaction


  1. Forward reaction

  2. Backward reaction
• Equilibrium


  The point where the forward and backward reactions are occurring at the same rate, so there is no overall change in concentrations
• Position of equilibrium


  The relative concentrations of reactants and products at equilibrium, which can shift left or right depending on conditions
• Adding heat to the reaction


  Encourages the forward reaction, shifting the position of equilibrium to the right
• Cooling the reaction


  Pushes the position of equilibrium back to the left
• Reversible reactions can only reach equilibrium in a closed system, where no reactants or products can escape</b>
• Exothermic


  A reaction that releases heat energy
• Endothermic


  A reaction that absorbs heat energy
• Reversible reactions are always exothermic in one direction and endothermic in the other
• Hydrated means water is present, anhydrous means no water is present
• Le Chatelier's Principle


  Principle about the position of equilibrium during a reversible reaction and how it's affected by temperature, pressure, and concentration
• Position of equilibrium


  Ratio of reactant particles to product particles when at equilibrium