1. Energy matters Unit 1

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 general to measure the rate of a reaction we must choose some measurable quantity which changes as the reaction proceeds. e.g mass of reactants in a flask, volume of gas produced, colour intensity, concentration of reagent

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. CaCO 3 (s) + 2HCl(aq) CaCl 2 (aq) + CO 2 (g) + H 2 O(l)

5. Marble chips & acid As we are producing a gas, it will escape from 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. Marble chips Balance Cotton wool HCl(aq)

6. Time (s)Total mass of flask (g) Decrease in mass (g) or Mass of CO 2 produced 0 30 60 90 120 150 180 210 240 270 300 149.00 147.75 147.08 146.60 146.26 145.94 145.68 145.48 145.32 145.19 145.08 - 1.25 1.92 2.40 2.76 3.06 3.32 3.52 3.68 3.81 3.92

7. Time (seconds) Decrease in mass (g)

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. Average reaction rate = Change in mass of product Time taken for change

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

10. 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.

11. 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.

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

13. 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.

14. 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. Activation energyThis minimum kinetic energy is called the Activation energy (E A )

15. 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.

16. Kinetic energy At a given temperature (T 1 ) 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.

17. 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 EAEA

18. Temperature Distribution of energy changes when the temperature changes. A small rise from T 1 to T 2 considerably increases the number of particles capable of reacting. Hence increasing the reaction rate. T2T2 Kinetic energy

19. 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.

20. 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.

21. 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.

22. 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. CO, NO x, O 2 CO 2 H 2 O O 2

23. Industrial catalyst CatalystProcessImportance Vanadium(v) oxideContactManufacture of H 2 SO 4 IronHaberManufacture of ammonia PlatinumOxidation of ammonia Manufacture of nitric acid NickelHydrogenationManufacture of margarine

24. 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.

25. Enzymes continued 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.

26. Industrial enzymes EnzymeProcess LipaseEnhance flavour of cheese, ice-cream & chocolate RenninCheese production ProteaseTenderising meat AmylaseDesizing (removing starch from fabric)

27. Enthalpy From SG: Exothermic reaction Combustion Neutralisation

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

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

30. Enthalpy enthalpy changeThe difference in potential energy between reactant and product is called the enthalpy change (ΔH) Enthalpy changes are normally quoted in kJ mol -1

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

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

33. Activated complex activated complexWhen reactants change into products, they pass through a very unstable state known as the activated complex. (Situated at the maximum potential energy). 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.

34. Activated complex

35. Patterns in the periodic table

36. 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.

37. 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.

38. 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.

39. 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.

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

41. 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) K 2+ (g)+e

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

43. 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. Screening / Shielding effect. Therefore the energy required to remove that electron decreases. Li Na K e-e- e-e- e-e-

44. Bonding, structure and properties of compounds

45.  Metallic Bonding  Covalent bonding  Polar covalent bonding  Ionic bonding

46. Metallic Lattice

48. Covalent molecular

49. Carbon atoms Covalent bonds Covalent network

50. Electronegativity The greater the difference in electronegativity between two elements, the less likely they are to share electrons and form covalent bonds. Caesium fluoride is the compound with the greatest degree of ionic bonding.

51. Formed when atoms of different electronegativities bond to form a covalent compound. Bonding electrons are not shared equally. The atom with the greater share of electrons becomes slightly negative (δ-) The other atom becomes slightly positive (δ+) These molecules have a permanent dipole. Polar covalent bonding

53. Ionic bonding Different elements have different attraction for bonding electrons, (electronegativity values). One atom may attract electrons very strongly and another atom may attract them very weekly and lets them go.

55. Ionic bonding

56. Summary

58. Intermolecular forces of attraction

59. Covalent molecular

60. Intermolecular interactions Van der Waal forces are a result of electrostatic attraction between temporary dipoles and induced dipoles caused by movement of electrons in atoms and molecules. All covalent molecules interact by van der Waals bonding, as all molecules possess temporary dipoles.

61. Halogens All halogen have 1 unpaired electron in the outer shell. Therefore form 1 pure covalent bond. E.g. F 2, Cl 2, Br 2, I 2 These molecules interact only weakly by van der Waals’ mechanism, this makes them very volatile. (Fluorine & chlorine are gaseous).

62. Permanent dipole A molecule can be described as polar if it has a permanent dipole. A permanent dipole is due to a difference in electronegativity between the atoms involved in a covalent bond.

64. Symmetry Some molecules have a symmetrical arrangement of polar bonds. This cancels out the polarity over the molecule as a whole.

66. Polar or Non-polar?

67. Boiling point Polar molecules have higher boiling points than non-polar molecules with similar molecular mass.

68. Hydrogen bonds Bonds consisting of a hydrogen atom bonded to an atom of a strongly electronegative element such as fluorine, oxygen or nitrogen.

69. Water molecules

70. Ice

72. GlycerolSulphuric Acid Phosphoric Acid

73. Covalent molecular

74. Carbon atoms Covalent bonds Covalent network

75. Diamond

76. Fullerenes Discrete covalently bonded molecules Consisting of pentagonal & hexagonal panels.

77. Graphite

78. Bonding, structure & properties of elements

79. Groups 1,2 & 3 Not enough electrons to achieve full outer shell. Elements contribute electrons to a common ‘pool’ of delocalised electrons. This binds the resultant positive ions. Bonding is less directional, therefore metals are more ductile & malleable. Delocalised electrons, therefore conduct electricity.

80. Metallic Bonding

81. 1 exception: Boron Structure made up of B 12 groups, interbonded with other groups. This results in an element almost as hard as diamond.

82. Group 4 Standard structure: Infinite 3D network or lattice, e.g. diamond, silicon. Therefore exceptionally hard & rigid. No discrete molecules, each atom joined to another.

83. Diamond

84. Graphite

85. Fullerenes Discrete covalently bonded molecules Consisting of pentagonal & hexagonal panels.

86. Phosphorus (group 5) Phosphorous bonds to 3 other phosphorous atoms to form tetrahedral P 4 molecules. Fewer electrons in P 4 than S 8 make van der Waals forces weaker in phosphorous, therefore lower m.p.

87. Group 6 Oxygen: 2 unpaired electrons, therefore forms 2 pure covalent bonds. Intermolecular interactions are weak van der Waals, therefore volatile & gaseous.

88. Sulphur Sulphur atoms can bond to more than one other sulphur, forming an 8 member ring. Van der Waals forces strong enough to make sulphur a solid at room temperature.

89. Groups 5, 6 & 7 Intra molecular forces (bonds within molecules) are covalent. Intermolecular forces are very weak van der Waals forces. Therefore most elements are volatile even if solid at room temperature. This is due to the little energy required to break intermolecular forces in order to melt/boil.

90. Bonding in elements: Noble gases 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.

91. Solvent action In general polar solvents dissolve polar substances and ionic substances.

92. Non polar solvents…(e.g hexane) Dissolve non polar solvents

93. The mole

94. The Avagadro constant 1 mole of any element contains the same number of atoms. This number is known as the Avagadro constant. This constant is given the symbol (L) after the first person to calculate a numerical value.

95. Avagadro constant (L) One mole of any substance contain L, 6.02x10 23 formula units.

96. Formula units For metals & monatomic species e.g. Noble gases, a formula unit is an atom. Thus 4g helium 40g of calcium 197g of gold Contain L (6.02x10 23 ) atoms

97. Covalent substances A formula unit is a molecule The total number of atoms can be found by multiplying L by the number of atoms in the molecule. Quantity of substance Number of molecules No. of atoms per molecule Total No. of atoms 2g of Hydrogen, H 2 18g of Water, H 2 O 30g of ethane, C 2 H 6 LLLLLL 238238 2L 3L 8L

98. Ionic compounds Formula unit consists of a ratio of ions expressed by ionic formula. Quantity of substanceNo. of formula units No. of +ve and –ve ions Total No. of ions 58.5g of Na + Cl - 74g of Ca 2+ (OH - ) 2 342g of (Al 3+ ) 2 (SO 4 2- ) 3 LLLLLL LNa + and LCl - LCa 2+ and 2LOH - 2L Al 3+ and 3L SO 4 2- 2L 3L 5L

99. Example 1 How many molecules are there in 8.8g of CO 2 ? 1 mole of CO 2 contains L molecules 44g of CO 2 contains L molecules 1g of CO 2 contains L/44 molecules 8.8g of CO 2 contains L/44 x 8.8 molecules = 1.204 X 10 23 molecules

100. Example 2 What mass of Nitrogen gas contains 18.06x10 22 atoms of Nitrogen? 6.02x10 23 molecules of N 2 1 mole 6.02x10 23 molecules of N 2 28g 1 molecule of N 2 28/L 18.06x10 22 molecules of N 2 28/L x 18.06x10 22 = 8.4g Therefore 8.4g of N 2 gas contains 18.06x10 22 molecules 4.2g of N 2 gas contains 18.06x10 22 atoms