C2 Rates and Energy

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  • Created by: Phoebe
  • Created on: 05-01-13 17:32

4.1 How fast?

The rate of a chemical reaction measures the speed of a reaction or how fast it is.

Rate of reaction =

amount of reactants used  OR  amount of product formed                            time                                       time

The gradient or slope of the line on a graph of amount of reactant or product against time tells us the rate of reaction at that time. The steeper the gradient, the faster the reaction.

The faster the rate, the shorter the time it takes for the reaction to finish. So rate is inversely proportional to time.

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4.2 Collision Theory and Surface Area

The Collision Theory states that reactions can only happen if the particles collide with enough energy to change into new substances.

The minimum amount of energy that particles must have in order to react is called the activation energy.

The rate of a chemical reaction increases if the surface area of any solid reactants is increased. By breaking a large solid into smaller fragments, we expose new surfaces, and this increases the frequency of the collisions.

Factors that increase the chance of collisions, or the energy of the particles, will increase the rate of reaction, e.g. temperature, concentration, pressure, surface area, use of a catalyst.

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4.3 The effect of Temperature

Reactions happen more quickly as the temperature increases.

Increasing the temperature increases the speed of the particles in a reaction mixture.

This means they collide more frequently and with more energy, which increases the rate of reaction.

At higher temperatures more of the collisions result in a reaction because a higher proportion of particles have energy greater than the activation energy.

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4.4 The effect of Concentration or Pressure

The particles in a solution are moving around randomly. If the concentration of a solution is increased there are more particles dissolved in the same volume. This means the dissolved particles are closer together and so they collide more frequently.

Increasing the concentration of a reactant therefore increases the rate of reaction because the particles collide more frequently.

In a similar way, increasing the pressure of a gas puts more molecules into the same volume, and so they collide more frequently. This increases the rate of reactions that have gases as reactants.

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4.5 The effect of Catalysts

A catalyst speeds up the rate of a chemical reaction.

They lower the activation energy of the reaction so that more of the collisions result in a reaction.

A catalyst is not used up in a chemical reaction.

Solid catalysts have large surface areas to make them as effective as possible.

Different catalysts are needed for different reactions.

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4.6 Catalysts in Action

Catalysts are used in industry to increase rate of reaction and reduce energy costs.

Traditional catalysts are often transition metals or their compounds which can be toxic and harm the environment if they escape.

Modern catalysts are being developed in industry which result in less waste and are safer for the environment.

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4.7 Exothermic and Endothermic reactions

Energy may be transferred to or from the reacting substances in a chemical reaction.

A reaction in which energy is transferred from the reacting substances to their surroundings is called an exothermic reaction, e.g. combustion, oxidation reactions and neutralisation reactions.

A reaction in which energy is transferred to the reacting substances from their surroundings is called an endothermic reaction, e.g. thermal decomposition of CaCO3

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4.8/9 Energy and Reversible reactions

In reversible reactions, the reaction in one direction is exothermic and in the other direction it is endothermic.

In any reversible reaction the amount of energy released when the reaction goes in one direction is exactly equal to the energy absorbed when the reaction goes in the opposite direction.

Exothermic changes can be used in hand warmers and self-heating cans.

Endothermic changes can be used in instant cold packs for sports injuries.

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