Fast and Slow Chemistry

1. Fast and Slow Chemistry Chapter 15

2. Fast and Slow Chemistry • Read page 247 • What would happen if this was a slow reaction???

3. Chemical Energy – What is chemical energy? • All substances have chemical energy. • The chemical energy of a substance is the sum of its potential energy (store energy) and kinetic energy (energy of movement).

4. What is chemical energy? • These energies result from such things as: • Attractions between electrons and protons • Repulsions between nuclei • Repulsions between electrons • Movement of electrons • Vibrations of and rotations around bonds • The chemical energy of a substance is sometimes called its heat content or enthalpy. • It is given the symbol H

5. Energy changes during chemical reactions • During a chemical reaction, the atoms in the reactants are arranged into products with different chemical energies.

6. Energy changes during chemical reactions – Exothermic • The total chemical energy of the products is less than the energy of the reactants. • Since energy is never lost, the difference in energy between reactants and products is released into the environment. • It is often released as heat energy.

7. Energy changes during chemical reactions – Endothermic • The chemical energy of the products is greater than the energy of the reactants. • Energy must be absorbed from the environment around the reactants in order for the reaction to occur.

8. Energy Changes • The energy released or absorbed during a chemical reaction is called the heat of reaction. • Since the heat of reaction is equal to the difference in enthalpy between the products and the reactants, it is given the symbol ∆H, where: • ∆H = H(products) – H(reactants)

9. ∆H • For exothermic reactions, ∆H will be negative • For endothermic reactions, ∆H will be positive • This depends on whether H(products) is greater than or less than H(reactants). • Most reactions we encounter are exothermic.

10. Thermochemical Equations • Thermochemical equations show the energy released or absorbed during a chemical reaction. • Energy is measured in Joules (J) or kilojoules (kJ). • 6CO2(g)+6H2O(l)→C6H12O6(aq)+6O2(g); ∆H=+2803 kJ mol-1 • C6H12O6(aq)+6O2(g)→ 6CO2(g)+6H2O(l); ∆H=-2803 kJ mol-1

11. Activation Energies CH4(g)+2O2→ CO2(g)+2H2O(g); ∆H=-890 kJ mol-1 • This is an exothermic reaction. • The energy of the reactants is higher than the energy of the products. • Why doesn’t natural gas burst immediately into flame and release energy when it comes into contact with air? • To start a gas oven, why must we use a match or a spark?

12. Activation Energies • Well what happens to chemical bonds during a reaction? • The bonds between the atoms in reactants must first be broken. • For this to occur energy must be absorbed • The new bonds form as the products are created • Energy is released as this happens

13. Activation Energy • The energy required to break the bonds of reactants is called the activation energy. • A diagram showing this is called an energy profile.

14. Your Turn • Pg 250 • Questions 1 and 2

15. Making Reactions Go Faster • The rate at which chemical reactions occur is an important consideration for industrial chemists and chemical engineers. • Some reactions take less than 10-11 seconds, while others take years. • Considerable effort is directed towards maximising reaction rates in industry.

16. 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’ • The rate of reaction can depend on the number of collisions as well as the energy of the collisions being greater than the activation energy.

17. Factors that affect rates • There are four main ways in which reaction rates can be increases: • Increasing the surface area of solids • Increasing the concentration of reactants in solution (or pressure of gases) • Increasing the temperature • Adding a catalyst • Explain to me how these factors can increase the rate of reaction?

18. Extending Collision Theory

19. Catalyst • Many reactions occur more rapidly in the presence of particular elements or compounds. • These substances, known as catalysts, are not consumed during the reactions and therefore do not appear as either reactants or products in reaction equations.

20. Catalysts • Many catalysts have been discovered by simple trial and error.

21. Catalysts • There are two types of catalysts • Homogenous catalysts – these are in the same state as the reactants and products • Heterogeneous catalysts – these are in different states from the reactants. • Chemists prefer to use Heterogeneous catalysts as they are more easily separated from the products of a reaction

22. How do catalysts work • Particles tend to adsorb to the surface of the catalyst. • Adsorption distorts bonds in the reactants allowing the reaction to proceed more easily than it would if the catalyst was absent. • Essentially a catalyst lowers the activation energy required to break the bonds of the reactants. • The relative energies of the reactants and products are unaffected by the presence of the catalyst. • This means ∆H is not changed.

23. How do catalysts work

24. Over to you • Read the Extension task on page 257 • Over to you • Page 258 • Question 3 and 4