1. Chapter 15 – Fast and Slow Chemistry

2. Fast and Slow Chemistry
During chemical reactions, particles collide and undergo change during which atoms are rearranged to produce new particles.
However, collisions between reactants do not always result in chemical change.

3. Chemical Energy All substances have chemical energy.
The chemical energy of a substance is the sum of its potential energy (stored energy) and kinetic energy (energy of movement).
The chemical energy of a substance is sometimes called heat content, or enthalpy, and is given the symbol H.
These energies can result from:
Attractions between electrons and protons
Repulsions between nuclei
Repulsions between electrons
Movement of electrons
Vibrations and rotations around bonds.

4. Exothermic Reaction
The total chemical energy of the products might be less than the energy of the reactants.
Since energy is never lost, the difference in energy between the reactants and products is released into the environment, usually as heat energy.
A reaction that releases energy is called an exothermic reaction.

5. Endothermic Reactions
The chemical energy of the products may be greater than the energy of the reactants.
Energy must be absorbed from the environment around the reactants in order for this reaction to occur.
A reaction that absorbs energy is called an endothermic reaction.

7. Endothermic and Exothermic Reactions
For exothermic reactions, H(products) is less than H(reactants), so values of ΔH will be negative – ΔH<0
For endothermic reactions, H(products) is greater than H(reactants), so values of ΔH will be positive – ΔH>0.

8. Thermochemical Equations
Photosynthesis is an endothermic reaction.
The thermochemical equation for this reactions is:
6CO2(g) + 6H2O(l)  C6H12O6(aq) + 6O2(g); ΔH = +2803kJ
In this reaction 2803kJ of energy is absorbed when 6 moles of carbon dioxide reacts completely with 6 moles of water to produce one mole of glucose and 6 moles of oxygen.
Thermochemical equations show the energy released or absorbed during a chemical reaction.
If the reverse reaction were to occur, the ΔH would be the opposite, i.e kJ.
If the equation were to be written for a reaction between 12 moles of carbon dioxide and 12 moles of water, then the ΔH would also be doubled.

9. Activation Energy
In an exothermic reaction, the energy of the reactants is higher than that of the products.
However, sometimes these reactions do not immediately occur.
Using the reaction between methane and oxygen to produce carbon dioxide and water, it can be explained why this is the case.

10. Activation Energy
When natural gas comes into contact with the air, why doesn’t it burst immediately? Why do we need a match to light a gas oven?
To answer these questions, we need to recall what happens to chemical bonds during the course of a reaction.
The bonds between atoms in the reactants must first be broken. This this to occur, energy must be absorbed.
Then new bonds form as the products are created. Energy is released when this happens.

11. Activation Energy
In the combustion of methane, energy is used to break covalent bonds within the methane and oxygen molecules.
Energy is then released when covalent bonds form as carbon dioxide and water are produced.
Since the combustion of methane reacts in an overall release of thermal energy, the energy absorbed when the bonds break must be less than the energy released when new bonds form.
The heat of a reaction is the net result of the energy absorbed in breaking bonds and the energy released by making them.
Energy changes are shown in a diagram called an energy profile.

12. Energy Profiles

13. Energy Profiles
The top of the curve represents an intermediate step in the reaction, referred to as the activation or transition complex.
At this stage the bonds are partially broken and the bonds in the products are partially formed.
The energy required to break the bonds of reactants so that a reaction can take place is called the activation energy.

14. Collision Theory
For a chemical reaction to occur, the particles involved must collide with each other with sufficient energy to overcome the activation energy ‘barrier’ to the reaction.
This way of visualising reactions is known as collision theory.
It is an extension of the kinetic theory of matter and covers the detail of what happens at the particle level during a chemical reaction.
The greater the number of collisions, the greater the rate of reaction.

15. Factors the Affect Rate
There are four main ways in which reaction rates can be increased:
Increasing the surface area of solids.
Increasing the concentration of reactants in solution (or pressure of gaseous reactants).
Increasing the temperature.
Adding a catalyst.

16. Increasing the Surface Area of Solids
In a solid, only those particles that are at the surface can be involved in reactions.
Crushing a solid into smaller parts means that more particles are present at the surface.
As a consequence of the greater number of exposed particles, the frequency of collisions between these particles and reactant particles increases, and so reaction occurs more rapidly.

17. Increasing the Concentration of Reactants
Reaction involving molecules or ions dissolved in solution occur faster if the concentration of the dissolved particles is increased.
With more particles moving randomly in a given volume of solution, the frequency of collisions is increased and so more successful collisions occur.
Increasing the pressure of gases raises the concentration of gas molecules, causing more frequent collisions.

18. Increasing the Temperature
As temperature increases, the average speed and average kinetic energy of the particle increases as well.
Therefore, as the temperature increases, so does the rate of reaction.

19. Extending the Collision Theory
The effect of temperature on reaction rate cannot simply be explained by the increased frequency of collisions.
A temperature increase of 10°C causes the rate of many reactions to double, but collisions have only increased by 1/50th of this amount.
At any given temperature, particles have a wide spread of kinetic energies and velocities.
Consequently, temperature is a measure of the average kinetic energy.

20. Extending the Collision Theory
If we increase the temperature of the particles, more particles will move at higher speeds and have higher kinetic energies.
There is still a wide spread of energies but a higher temperature will mean more particles have higher kinetic energies.

21. Extending the Collision Theory
A reaction can take place only if the particles colliding have more energy than the activation energy of the reaction.
At any instant only a fraction of the particles present have sufficient energy to participate in successful collisions and react.
EA represents the activation energy of a reaction.

22. Catalysts
Many reactions occur more rapidly in the presence of particular elements or compounds.
These substances, catalysts, are not consumed during the reactions and therefore do not appear as either reactants or products.
There are two types of catalysts:
Homogenous catalysts: are those in the same state as the reactants and products.
Heterogeneous catalysts: are in different states from the reactants.

23. How do catalysts work?
In general, particles at the surface of a solid of high surface energy tend to adsorb, for a bond with, gas molecules that strike the surface, lowering the surface energy of the solid.
Adsorption distorts bonds in the gas molecules or may even break them completely, allowing a reaction to proceed more easily than it would if the solid were absent.
A powdered or sponge-like form of the solid catalyst is used to provide the greatest possible surface area.
The larger the surface area, the more reactant molecules that can be adsorbed, and the faster the reaction.

24. How do catalysts work

25. How do catalysts work
The catalyst provides an alternative reaction pathway that dramatically reduces the activation energy.
Consequently, a much higher proportion of reactant particles collide with sufficient energy to overcome the activation energy barrier.
Providing this pathway increases the proportion of successful collisions and therefore the reaction rate is increased.
In the diagram on the previous slide, the energies of the reactants and products do not change, meaning ΔH does not change.
Adding a catalyst only changes the activation energy of a reaction.