1. Chemical Energy Equilibrium

2. Chemical Energy The chemical energy of a substance is the sum of its potential energy (stored energy) and kinetic energy (energy of movement). 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

3.  All chemical reactions result in a change in the chemical energy of the reactants as they form the products Reaction Energy

4. During a chemical reaction, the atoms in the reactants are arranged into products with different amounts of chemical energies.

5. Reaction Energy Chemical changes that give off energy are called exothermic reactions. Chemical changes that absorb energy are called endothermic reactions.

6. Exothermic The total chemical energy of the products is less than the energy of the reactants. the difference in energy is released into the environment, often as heat energy.

7. Exothermic Reactions 25 o C 45 o C Gets hot Hydrochloric acid magnesium Magnesium + Hydrochloric acid

8. Energy Level Diagram: Exothermic Reaction Energy / kJ) Progress of reaction (time) reactants Reactants have more chemical energy. Some of this is lost as heat which spreads out into the room. products Products now have less chemical energy than reactants.

9. Exothermic Reaction Rate Diagram

10. Endothermic The chemical energy of the products is greater than the energy of the reactants. Energy must be absorbed from the surrounds in order for the reaction to occur.

11. Endothermic Reactions Ammonium nitrate Water Cools Heat energy taken in as the mixture returns back to room temp.

12. products Energy / kJ) Progress of reaction reactants HH This is how much energy is taken in Energy Level Diagram: Endothermic Process

13. Reaction Energy Why doesn’t natural gas burst immediately into flame when it comes into contact with air? Why must matches be struck for them to light? To answer these questions you need to think about what happens to the chemical bonds of the substances during a reaction.

14. Reaction Energy The bonds between the atoms in the reactants must first be broken. – For this to occur energy must always be absorbed New bonds form as the products are created – Energy is always released as this happens

15. Reaction Energy Step 1: Energy must be SUPPLIED to break bonds: Step 2: Energy is RELEASED when new bonds are made: A reaction is EXOTHERMIC if more energy is RELEASED then SUPPLIED. If more energy is SUPPLIED then is RELEASED then the reaction is ENDOTHERMIC

16. Burning Methane CH 4 + 2O 2 2H 2 O + CO 2 These bonds must be broken : These bonds must form

17. 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.

18. Why is Switzerland like a chemical reaction?

19. Activation Energy Chemical Reactions must go over an energy hill like a car over the mountains (Swiss Alps).

20. Collision Theory For a chemical reaction to occur, the particles involved must collide with each other. The collisions must be with sufficient energy to overcome the activation energy ‘barrier’. The rate of reaction (how quickly the reaction occurs) depends on the number of energy sufficient collisions per time.

21. Collision Model Molecules must collide in order for a reaction to occur. What factors affect the rate of collisions?

22. Factors That Affect Reaction Rates There are four main ways in which reaction rates can be increases: – Increase the surface area of solid reactants (crush) – Increase the concentration of the reactants – Increasing the temperature of the reactants – Adding a catalyst Explain how each of these factors can increase the rate of reaction.

23. Surface Area Only the reactants that are on the surface of a cube are able to react. They are the only molecules exposed to other reactants. The greater the number of exposed molecules the larger the number of possible collisions. Take a cube 2 X 2 X 2 cm. Find the total surface area. Brake the cube into 8 separate cubes of 1 X 1 X 1 cm. Find the total surface area.

24. Concentration As the concentration of the reactants goes up the distance between the molecules becomes less. With less distance to travel the time between collisions becomes less

25. Temperature The faster the molecules move the less time it take for them to reach another molecule

26. 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, they just help it along They often work by lowering the activation energy. Enzymes are biological catalysts.

27. Catalyst

28. Catalytic converters are used to convert the polluting exhaust gases of burned lead-free gasoline into harmless gases. Platinum (Pt) is the catalysts used. Only a small amount is needed.

29. 7.5 Equilibrium Some go to completion All the reactants get turned into products No reactants left Some reactions go both directions They are called reversible reactions

30. Chemical Equilibrium Eventually you reach a point where the reverse reaction is going as fast as the forward reaction. This is chemical equilibrium. The rate of the forward reaction is equal to the rate of the reverse reaction. The concentration of products and reactants stays the same, but the reactions are still running.

31. Chemical Equilibrium A dynamic state where the concentrations of all reactants and products remain constant.

32. Chemical Equilibrium In a chemical equilibrium, forward and reverse reactions occur at the same rate; the concentration of the reactants and products remain constant. You can cause the equilibrium to shift by adding or removing reactants or products. Below CO 2 gas is in equilibrium with the water in the soda.

33. Equilibrium Equilibrium position- how much product and reactant there are at equilibrium. Shown with the double arrow.

34. Reversible Reactions 2H 2 (g) + O 2 (g)  2H 2 O(g) + energy 2H 2 O(g) + energy   H 2 (g) + O 2 (g) 2H 2 (g) + O 2 (g)  2H 2 O(g) + energy

35. LeChâtelier’s Principle When a change is introduced to a system in equilibrium, the equilibrium shifts in the direction that relieves the change. Three types of changes are described. – Temperature – Pressure – Concentration N 2 + 3H 2 2NH 3 + Heat

36. How ammonia is produced N 2 + 3H 2 2NH 3 + Heat Temperature  If heat is added, the equilibrium will shift to remove heat from the system (since heat is a product it would lower amount of NH3) Pressure  If pressure is increased, will shift to decrease pressure (fewer molecules on products, increase NH3) Concentration  change in amount of molecules shift reaction towards lowered amount

37. So at low temperatures, high pressure and low concentration of ammonia This maximizes the amount of ammonia produced.

38. Changing Temperature Reactions either require or release heat. Endothermic reactions go faster at higher temperature. Exothermic go faster at lower temperatures. All reversible reactions will be exothermic one way and endothermic the other.

39. Changing Temperature As you raise the temperature the reaction proceeds in the endothermic direction. As you lower the temperature the reaction proceeds in the exothermic direction.

40. Changing Concentration If you add reactants (or increase their concentration). The forward reaction will speed up. More product will form. Equilibrium “Shifts to the right” Reactants  products N 2 + 3H 2 2NH 3 + Heat

41. Changing Concentration If you add products (or increase their concentration). The reverse reaction will speed up. More reactant will form. Equilibrium “Shifts to the left” Reactants  products

42. Changing Concentration If you remove products (or decrease their concentration). The reverse reaction will slow down More product will form. Equilibrium reverse“Shifts to the right” Reactants  products

43. Changes in Pressure As the pressure increases the reaction will shift in the direction of the least gases. At high pressure 2H 2 (g) + O 2 (g)  2 H 2 O(g) At low pressure 2H 2 (g) + O 2 (g)  2 H 2 O(g)