Energy matters

1. Energy matters Higher chemistry

2. Reaction rates • From standard grade you should remember that a reaction can be speeded up by; • Decreasing particle size • Increasing concentration • Increasing temperature

3. Following the course of a reaction • In some reactions we can monitor the rate at which the reaction is taking place. • We do this by taking measurements over a period of time.

4. Following the course of a reaction • If we react marble chips (Calcium carbonate) with hydrochloric acid we can monitor the course of the reaction. CaCO3(s) + 2HCl(aq) CaCl2(aq) + CO2(g) + H2O(l)

5. Marble chips & acid • As we are producing a gas, it will escape for the vessel causing the total mass to drop. • If we measure this change in mass over a fixed period of time we can calculate the rate of the reaction. Cotton wool Marble chips HCl(aq) Balance

8. Average rate of reaction • It is difficult to measure the actual rate at any one instant since the rate is always changing. • We can calculate average rate over a certain period of time. Change in mass of product Average reaction rate = Time taken for change

9. Example • Calculate average rate of reaction between 30 and 60 seconds. Change in mass of product Average reaction rate = Time taken for change 1.92 - 1.25 Average reaction rate = 30 0.022gs-1 Average reaction rate =

10. Work • Finish PPA experiments (1&2) • PPA write ups (1&2) • Rates worksheets. • Rates homework sheets.

11. Collision theory • For a chemical reaction to occur, the reactants must collide. • Any factor that increases the number of collisions per second is likely to increase reaction rate.

12. Particle size • More collisions occur if the particle size of a solid reactant is decreased, since its overall surface area is increased. • Powdered marble (calcium carbonate) reacts much faster than marble chips.

13. Concentration • If concentration is increased, there are more reactant particles. • The more particles there are in one space, the more collisions.

14. Raising temperature • Raising the temperature at which a reaction takes place does more than merely raise the number of collisions. • Temperature is a measure of the average kinetic energy of particles in a substance. • Therefore at higher temperatures, particles have greater kinetic energy and they collide with more force.

15. Collisions • Not all collisions cause a reaction to occur e.g. nitrogen & oxygen particles in the air. • The colliding particles must have a minimum amount of kinetic energy for a reaction to occur. • This minimum kinetic energy is called the Activation energy (EA)

16. Activation energy • Activation energy required varies from one reaction to another. • If the activation energy of a reaction is high, only few particles will have enough energy to successfully collide. • Conversely, a reaction with low activation energy will be very fast.

17. Kinetic energy • At a given temperature (T1) individual molecules of a gas have widely different kinetic energies. • Most molecules will have energy near to the average energy but some will be well below average, and some well above.

18. Activation energy • The shaded area represents the all of the molecules which have kinetic energy greater than the activation energy. • The shaded area represents the portion of molecules that will react EA

19. Temperature • Distribution of energy changes when the temperature changes. • A small rise from T1 to T2 considerablyincreases the number of particles capable of reacting. • Hence increasing the reaction rate. T2 Energy (kJ)

20. Catalysts • Substance that alters rate of reaction without being used up. • Homogeneous catalyst: Same state as the reactants. • Heterogeneous catalyst: Different state as the reactants.

21. Transition metals • Are characterised by: • Catalytic behaviour • Variable valency • Formation of coloured ions.

22. Industrial catalyst

23. Heterogeneous catalyst • The catalyst has a large surface area. • Catalysis occurs at certain points on the catalyst called ‘active sites’. • At these sites reactant molecules are adsorbed onto the surface of the catalyst. • At least 1 reactant is held in place on active site, making collision more likely. Copy diagram from higher chemistry text book page 11

24. Catalyst poisoning • Occurs when reactants or impurities become preferentially adsorbed or even permanently attached to the catalyst surface. • Hence reducing number of active sites and therefore rendering the catalyst as useless.

25. Catalytic converters • Petrol engine cars must now be fitted with a catalytic converter. • The contains a honeycomb network of platinum, converting harmful gases into less harmful ones. CO2 H2O O2 Less harmful Harmful CO, NOx, O2

26. Enzymes • Biological catalyst. • Examples of enzymes: • Amylase, catalyses the hydrolysis of starch. • Catalase, catalyses the decomposition of hydrogen peroxide. Catalyase is found in the blood, preventing build up of hydrogen peroxide in the body.

27. Enzymes are highly specific. Enzymes work best at their optimum temperature & pH. Optimum temperature for human enzymes will be 37°C. Greatly exceeding either of these will result in the protein being denatured. Enzymes continued

28. Industrial enzymes

29. Enthalpy From SG: Exothermic reaction Combustion Neutralisation

30. Potential energy • Potential energy is the energy possessed by the reactants. • In an exothermic reaction, the products have less potential energy than reactants.

31. In an endothermic reaction, the opposite is true. Reactants must absorb energy from their surroundings. Products have more energy than the reactants. Potential energy

32. Enthalpy • The difference in potential energy between reactants is called the enthalpy change ( H) • Enthalpy changes are normally quoted in kJ mol-1

33. For an exothermic reaction, H is a negative value. e.g. H2(g) + ½ O2(g) H2O(l) H = -286kJmol-1

34. For an endothermic reaction, H is a positive value. e.g. C(s) + H2O(g) CO(g) + H2(g) H = +121 kJmol-1

35. The rate of reaction depends on the height of the Ea barrier. Rate of reaction does not depend on the enthalpy change ( ) Activation energy H H

36. Activated complex • When reactants change into products, they pass through a very unstable state known as the activated complex. • The activated complex is a highly energetic arrangement of atoms that exists for a short time. • The activated complex loses this energy by either forming products or reforming as reactant particles.

37. Activated complex

38. Catalyst • Catalysts provide alternative reaction pathways. • Thus lowering the activation energy. Energy Reaction pathway

39. Patterns in the periodic table

40. Density • The amount of material packed into a given volume. • Density values are much larger for Solid & liquid elements. • Density increases down each group. • Across the period from L to R, density increases towards the centre of the period, then decreases again towards the noble gases.

41. Atomic size: Groups • Atomic size is measured in covalent radius. This is the distance from the nucleus to the outer electrons. • As you move down a group the atomic radius increases. • This is due to the increased number of occupied electron shells. Increasing atomic radius

42. Atomic radius: Periods • Across a period atomic number and electron number increase by one. • Although the number of outer electrons is increasing across the period, the atomic radius decreases. • This is due to the increasing attraction between the nucleus and the outermost electrons. Decreasing atomic radius

43. Ionisation energies • The attraction between the nucleus and the outer electrons means that energy is required to remove electrons from the atom. • Ionisation energy is a measure of the nuclear attraction for outer electrons.

44. First ionisation energy • Energy required to remove an electron from one mole of free atoms in a gaseous state. • K(g) K+(g) + e-

45. Second ionisation energy • Energy required to remove an electron from one mole of ions with a charge of 1+ in the gaseous state. • K+(g) K2+(g) + 2e

46. Third ionisation energy • Energy required to remove an electron from 1 mole of ions with 2+ charge in the gaseous state. • K2+ K3+(g) + 3e

47. Ionisation energies • The first ionisation energy decreases as you go down a group. • This is due to the increasing atomic radius. • As the radius increases, the attraction between the nucleus and the outermost electrons decreases. • Therefore the energy required to remove that electron decreases. Li e- Na e- K e-

48. Bonding, structure & properties of elements

49. There are no covalent or ionic bonds between atoms in group 8. Uneven distribution of electrons within the atom produce temporary (or transient) dipoles on the atom. Bonding in elements: Noble gases van der Waals' forces

50. Halogens • All halogen have 1 unpaired electron in the outer shell. Therefore form 1 pure covalent bond. E.g. F2, Cl2, Br2, I2 • These molecules interact only weakly by van der Waals’ mechanism, this makes them very volatile. (Fluorine & chlorine are gaseous).