rate of reaction

• quantity of product formed (g)or (cm3)(mol)(dm3)

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  time taken (s)
• quantity of reactant used (g)(cm3)(mol)(dm3)

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  time taken (s)
• reaction is fastest in the beginning, slows down as it progresses.
• Identifying chemical changes:

  • Gas is made
  • Precipitate is made
  • Colour change
  • Change in temperature
• Rate of reaction: The speed at which the reactants are used up or at which new products are formed
• A catalyst increases the rate of a reaction and is unchanged at the end of a reaction
• Collision Theory:

  1. Number of particles per unit volume
  2. Frequency of collision between particles
  3. Kinetic energy of particles
  4. Activation energy
• • The SLOPE shows the rate of reaction
  • steeper=faster
• 1. Initially, the reaction is very fast (a lot of product, short amount of time). This is because there is a large number of reactant molecules, so more of them are reacting and forming a product.

  
• 2. Gradually, the slope becomes less steep. The reaction is slowing down and the rate of reaction is decreasing. This is because a lot of the reactant molecules have already reacted and formed a product. So, there are fewer reactant molecules available to react. (conc is decreasing & fewer collisions)

  
• 3. Slope = zero . The line is flat. The reaction has stopped because all the reactant molecules have reacted.

  
• 1. another graph

  A) initially fast

  B) slows down

  C) stops

  
• Use a tangent to find the mean rate of reaction

  
• Collision Theory: Chemical reactions can only take place when reacting particles collide with sufficient energy
• Frequency: The number of successful collisions per second. This is what determines the rate of reaction.
• more concentration will result in more product because the reaction started with more reactant molecules.
• More concentrated = More amount of solute than amount of solvent  per unit volume = Increase the frequency of collision = faster successful reactions = faster rate of reaction

  
• Larger surface area = more reactant particles exposed to the reaction = Increase the frequency of collision = speed up rate of reaction
• The pressure of a gas depends on the number of molecules. more molecules = higher pressure = more successful collisions = higher rate of reaction .
• activation energy: the minimum amount of energy needed for a reaction to occur / for particles to collide successfully
• increasing temp = increasing energy = increasing speed of particles = increases frequency of collisions

  AND more particles have sufficient energy (above activation energy) and can overcome the activation energy barrier and collide successfully.
• catalyst decreases the activation energy of a reaction. this saves money because they are reusable.
• catalysts provide an alternative pathway for the reaction that has a lower activation energy. more particles can successfully collide per second, which increases the rate of reaction

  
• DISAPPEARING CROSS:

  1. Use a measuring cylinder to pour 10cm3 of sodium thiosulphate in a conical flask.
  2. Place the conical flask on top of a printed black cross
  3. Add 10cm3 of hcl into the conical flask
  4. swirl the solution and start a stopwatch
  5. look down through the top of the flask, the solution should turn cloudy after some time.
  6. Stop the clock when you can no longer see the cross
  7. Redo the experiment using different concentrations of sodium thiosulphate solution
  8. Repeat the whole experiment and calculate means

  [Different people have different eyesights so it may not be as accurate]
• GAS PRACTICAL

  1. Use a measuring cylinder to pour 50cm3 of HCl in a conical flask
  2. Attach the conical flask to a bung and a delivery tube
  3. Place the delivery tube in a container full of water
  4. Place an upturned measuring cylinder also filled with water over the delivery tube.
  5. Add a 3cm strip of Mg in the HCl and start a stopwatch
  6. Every ten seconds, measuring the volume of Hydrogen gas trapped in the measuring cylinder until no more hydrogen is given off.
  7. repeat the experiment with different concentrations of HCl
• flour & coal

  • fine carbon reacts quickly with the oxygen in the air and cause a highly exothermic reaction