1. C2 Additional Chemistry

2. Ionic Bonding

3. Metals Non-Metals H have too many electrons don’t have enough Atoms want to have a full outer shell of electrons to become stable just like the noble gases

4. Metal atoms lose electrons to become positive ions. Non-metals atoms gain electrons to become negative ions. Oppositely charged ions have a strong electrostatic attraction towards each other. Li X X F X +1 Li X X F X Ionic Bonding

5. +1+2+2+3 Group 1 (and H + ) NH 4 + (Ammonium) Group 2Aluminium -2-2-3 Group 7 OH - (hydroxide) NO 3 - (nitrate) Group 6 CO 3 2- (carbonate) SO 4 2- (sulphate) PO 4 3- (phosphate) After ionic bonding we sometimes get differently charged ions...but compounds are always neutral! So to find the formula combine ions and make sure that the charges cancel each other out. E.g. Na 2 O, MgCl 2, Al(OH) 3, Al 2 O 3

6. Ionic Properties Oppositely charged ions are attracted to each other to form a giant ionic lattice. A lot of heat energy is needed to break the strong bonds so ionic compounds have high melting points. Ionic compounds can also conduct electricity when they are molten or in a solution because ions are free to move.

7. Metallic Bonding

8. Bonding & Structure 1.Metals lose their outer electrons to become positively charged ions 2.These ions form layers to give a giant metallic lattice 3.These are surrounded by a sea of delocalised electrons 4.The delocalised electrons are free to move Before Bonding After Bonding

9.  High melting points due to strong attraction between positive ions and electrons  Conduct electricity and heat because the electrons are free to move ++ + + + ++ + + - - - - - - - - - - Delocalised e - s are free to move Properties of Metals

10. Pure metals:  One type of atom only  Regular layers  Layers can slide easily  Malleable (soft) Alloys:  Mixture of metals  Distorted layers  Cannot slide easily  Much harder Pure Metal Alloy

11. Covalent Bonding

12.  Non-metal atoms only  Atoms share electrons  A molecule is formed

13. Simple covalent molecules consist of a set number of atoms. e.g. H 2 O, CO 2, O 2 They have weak intermolecular bonds between the molecules. Little energy is needed to break these bonds so they have low boiling points. C O O COO Weak bonds Simple molecules

14. Giant covalent structures consist of infinite atoms. E.g. Sand is made from silicon dioxide Strong covalent bonds  very high melting points Carbon atoms exists as 3 different giant structures Giant Structures Fullerene Silicon dioxide Diamond Graphite

15. Diamond:  Each carbon atom forms 4 covalent bonds  Strong bonds hold atoms in a rigid, pyramid structure  Extremely hard – used for cutting tools

16. Graphite:  Each carbon forms 3 covalent bonds  Flat, layered structure  Weak intermolecular forces between layers  Layers can slide over each other - soft & slippery  Delocalised electrons are free to move - conducts electricity

17. Properties of Polymers

18. Polymers are long molecules that are formed when lots of small monomers are joined together by covalent bonds. The bonds/forces between the polymer chains determines their properties = = = = - - - - - - - - - - - - - - - - - - - -

19. LDPE: Low density poly(ethene) HDPE: High density poly(ethene) Conditions High pressure, trace oxygen Catalyst, 50 o C Diagram PropertiesFlexible but weakStrong but brittle Types of poly(ethene)

20. Thermosetting plastics do not become soft when heated. They have strong covalent cross links between the polymer chains. These bonds do not break, keeping the chains together. Thermosoftening plastics become soft & melt when heated. They have weak forces between the polymer chains. Forces are easily broken allowing the chains to slide over each other. Thermosoftening vs Thermosetting

21. Nanoscience

22. Nanoscience is the study of tiny nanoparticles Nanoparticles contain just a few hundred atoms and are less than 100nm in size These have an extremely high surface area to volume ratio and have special properties and uses:  Cosmetics and suncream  Electronics  Catalysts Gold nanoparticles used in cancer therapy Nanotubes used in electronic products

23. Relative Mass

24. Li 3 7 P P P N N N N neutrons = mass – atomic number (top – bottom number) Mass of atoms The top number on the periodic table tells us the mass of an atom It tells us the number of protons + neutrons. (electrons have no mass)

25. Isotopes Isotopes are forms of the same element which have a different relative atomic mass. They have the same number of protons but a different number of neutrons. Cl 17 35 Cl 17 37 Protons = 17 Neutrons = 18 Protons = 17 Neutrons = 20

26. The relative atomic mass, A r, of an element is the average mass of its atoms compared to the isotope 12 C. Cu 29 63.5 Cl 17 35.5 O 8 15.99 Ar is often rounded to the nearest whole number. Relative Atomic Mass

27. 1.MgO = 24 + 16 = 40 2.H 2 O = 1 + 1 + 16 = 18 3.CO 2 = 12 + 16 + 16 = 44 Relative formula mass (M r ) = Sum of the relative atomic masses in a compound Mg 12 24 H 1 1 C 6 12 O 8 16 Relative Formula Mass

28. Percentage Mass & Yield

29. % mass = Ar x number of atoms x 100 Mr of compound relative atomic mass of element in question Small number after element in question Relative molecular mass of whole compound E.g. Find the % mass of O in CO 2 % mass = 16 x 2 x 100 12+16+16 = 72.7% C 6 12 O 8 16 Percentage Mass

30. Percentage Yield We do not always get as much product from a reaction as we expect. This may be because some of the product is left on the apparatus or spilt. % yield = Actual mass x 100 Expected mass E.g. 6g of ammonia was produced even though 24g was expected to be produced % mass = (6÷24) x 100 = 25% Reversible reactions give very low yields!!

31. Empirical Formulae

32. The empirical formula tells you the number of each type of atom in a compound. to find... 1.Show formula as a ratio 2.Divide the masses, or percentages, by the relative atomic masses (Ar) 3.Simplify ratio by dividing all answers by smallest number 4.Use numbers to write empirical formula

33. Examples a)A compound contains 64g of sulphur and 64g of oxygen. What is it’s empirical formula? 1.S : O 2.64 ÷ 32: 64 ÷ 16 2: 4 3.2 ÷ 2: 4 ÷ 2 1: 2 4.SO 2 S 16 32 O 8 16

34. Examples b)A compound contains 40% Ca, 12% C and 48% O. What is it’s empirical formula? Relative atomic masses,Ar : C=12, O=16, Ca=40 1.Ca : C:O 2.40 ÷ 40: 12 ÷ 12:48 ÷ 16 1: 1:3 3.Already simplified! 4.CaCO 3

35. Mole Calculations

36. Moles tells us how many particles are in a substance. In 1 mole there are 6x10 23 particles. A mole is basically just a really big number, sometimes known as Avogadro’s number E.g. Find the mass in grams of 2 moles of Na Mass= moles x Ar = 2 x 23 = 46g Na 11 23 A r or M r Moles Mass (g)

37. The number in front of each element or compound in a balanced equation gives us the mole ratio. 2Mg+ O 2  2MgO Mole ratio 2: 1 : 2 Mass (g) Some of these will be in the Q Relative mass Calculate but ignore mole ratio (ignore big numbers) Moles Calculate using triangle or ratio Mole ratios can be used together with the moles triangle to calculate the number of moles & masses of all reactants & products in a reaction. Remember to lay your answer out in this way...

38. 1. What mass of water is produced when 20g of hydrogen reacts with oxygen? Relative atomic masses (Ar): H=1, O=16 2H 2 +O 2  2H 2 O Mole ratio 2: 1: 2 Mass (g)20 180 Relative mass 2 32 18 Moles10 5 10 Found using mole ratio A r or M r Moles Mass (g) x

39. 2. What mass of water is needed to make 49g of sulphuric acid? Relative atomic masses (Ar): H=1, O=16, S=32 SO 3 +H 2 O  H 2 SO 4 Mole ratio 1: 1: 1 Mass (g) 9 49 Relative mass80 18 98 Moles0.5 0.5 0.5 Found using mole ratio A r or M r Moles Mass (g) x

40. Relative atomic mass (A r ): Top number (periodic table) Relative formula mass (M r ): Sum of atomic masses Percentage mass = Ar x number of atoms x 100 Mr of compound Ar or Mr Moles Mass (g) Empirical Formula: 1) Show formula as a ratio 2) Divide percentages by the Ar 3) Simplify ratio 4) Write empirical formula % yield = Actual mass x 100 Possible mass Summary of Calculations

41. Analysing Substances

42. Paper chromatography can be used to separate mixtures. The number of spots shows the number of substances. The height of the spots can be used to identify the substance Most Soluble Least Soluble Paper Chromatography

43. The mix consists of dyes 1 and 4 1 2 3 4 mix Using Paper Chromatography

44. GC-MS machines can be used to analyse mixtures of compounds. They are used to test for drugs in an athlete’s urine. Gas chromatography (GC) separates the mixture Mass spectrometry (MS) measures the relative molecular mass of the compounds. C 9 H 13 NO M r = 151 Analysing Mixtures Using GCMS

45. Gas chromatography (GC) 1.Sample is injected into a gas stream 2.Gas stream carries sample through a long thin column 3.Some compounds take longer to pass through the column

46. Mass spectrometry (MS) Molecular ion peak shows the relative molecular mass

47. Advantages of GCMSDisadvantages of GCMS Highly accurate and sensitive Quick results Can analyse very small samples Expensive Specialist training required Pros and Cons of GCMS

48. Speeding up reactions

49. Collision theory Particles need to collide with enough energy for a reaction to take place. This is called the activation energy. ABABC

50. Increase temperature Particles have more kinetic energy Particles collide more often and with more energy More collisions above the activation energy Rate of reaction increases Temperature

51. Smaller pieces of solid Larger surface area to volume ratio More collisions take place above the activation energy Rate of reaction increases Surface Area : Volume

52. Increase concentration/pressure More particles so the particles are more crowded Particles collide more often More collisions above the activation energy Rate of reaction increases Low pressure/ concentration High pressure/ concentration Concentration or Pressure

53. Catalysts speed up the reactions by lowering the activation energy. They remain unchanged at the end of the reaction so can be used over and over again. Different catalysts are used for different reactions. Transition metals are very good catalysts. without a catalyst with a catalyst -- activation energy reactants products Catalysts

54. Catalysts speed up many useful reactions The very expensive metal platinum is used in catalytic convertors: 2CO + O 2  2CO 2 Catalysts eventually get blocked up and have to be thrown away. This leads to toxic waste.

55. Enzymes are biological catalysts that can be used to speed up cleaning clothes  Lower temperatures can be used - saves money Nano-sized catalysts have a huge surface area to volume ratio  less catalyst is needed which saves money

56. Rates Summary Surface AreaTemperature Concentration/ Pressure Catalysts Smaller pieces (Powder) More collisions Faster particles More energy More collisions Lower activation energy Do not get used up Transition metals used More collisions More particles

57. Measuring Rates

58. Rates of reactions are measured by timing how quickly:  the mixture gets cloudy if a solid is formed  the mass decreases  gas is produced using a gas syringe (s) = solid (l) = liquid (g) = gas (aq) = aqueous Measuring Rates of Reaction

59. Rate = amount of product time taken Calculating the Rate Graph curves because:  Reactants get used up  Fewer collisions  Rate slows down  Reaction stops when reactants are used up The steeper the gradient, the faster the rate of reaction. fast slow Reaction stops Mass of products

60. Heat Changes

61. reactants products Energy heat Exothermic Reactions heat Exothermic – Heat energy is given out to the surroundings (temperature rises)

62. reactants products Energy heat Endothermic Reactions heat Endothermic – Heat energy is taken in from the surroundings (temperature drops)

63. Reversible Reactions

64. Forward Reaction Reverse Reaction Water escapes Water added Reversible reactions can go backwards Reactants Products

65. If the forward reaction is exothermic......the reverse reaction must be endothermic Reversible Heat Changes Exactly the same amount of energy will be released or absorbed

66. Hydrated Copper Anhydrous Copper + Water Sulphate Sulphate Example – Copper Sulphate Forward Reaction Reverse Reaction Heat IN Heat OUT ENDO- THERMIC EXO- THERMIC

67. What is Electrolysis?

68. Ionic compounds must be molten (liquid) or in solution (aqueous) so that the ions are free to move and conduct electricity. Electrolysis is used to break up/separate ionic compounds using electricity. E.g. Electrolysis of copper chloride: CuCl 2 (aq)  Cu(s) + Cl 2 (g)

69. + - + Anode(+) - Cathode(-) + - - + - + Positive ions move to the negative electrode Negative ions move to the positive electrode - +

70. Oxidation and Reduction

71. Oxidation Is Loss of electrons Reduction Is Gain of electrons OIL RIG

72. Cu 2+ Cl - + - Copper chloride, CuCl 2 (aq) Positive electrode: 2Cl - - 2e -  Cl 2 OXIDATION Negative electrode: Cu 2+ + 2e -  Cu REDUCTION 2e - e-e- e-e- Cu Cl Oxidation can be written as 2Cl -  Cl 2 + 2e - Half Equations

73. Negative electrode: 2H + + 2e -  H 2 REDUCTION H+H+ O 2- Water, H 2 O (l) Positive electrode: 2O 2- - 4e -  O 2 OXIDATION e - 2e - H O O Half Equations + - Oxidation can be written as 2O 2-  O 2 + 4e - H+H+ e - H

74. Extracting Aluminium

75. 2.Oxygen forms at the positive electrode and aluminium forms at the negative electrode. 3.The carbon electrodes must be replaced regularly because they react with the oxygen: C + O 2  CO 2 1.Aluminium oxide is mixed with molten cryolite.This lowers the melting point and reduces the energy needed to melt the ore. Aluminium is extracted from an ore called bauxite (aluminium oxide) by electrolysis.

77. Al 3+ O 2- + - Aluminium Oxide, Al 2 O 3 (l) Positive electrode: 2O 2- - 4e -  O 2 OXIDATION Negative electrode: Al 3+ + 3e -  Al REDUCTION 3e - 2e - Al O O Half Equations Oxidation can be written as 2O 2-  O 2 + 4e -

78. Electrolysis of Brine

79. +- + - Cl 2 H2H2 NaOH 3 useful products are made by the electrolysis of sodium chloride solution (NaCl): Positive electrode  Chlorine gas used for bleach, water treatment & plastics Negative electrode  Hydrogen gas used to make margarine Left in solution  Sodium hydroxide used to make soap, paper & bleach Why do we electrolyse brine?

80. + - + Anode(+) - Cathode(-) Cl - Na + Cl - Na + Cl - Na + Cl - NaCl (aq) contains Na +, Cl -, H + & OH - ions Cl - ions go to the positive electrode to make Cl 2 H + ions go to the negative electrode to make H 2 H+H+ H+H+ H+H+ H+H+ OH -

81. Negative electrode: 2H + + 2e -  H 2 REDUCTION H+H+ Cl - Brine, NaCl (aq) Positive electrode: 2Cl - - 2e -  Cl 2 OXIDATION e - e-e- e-e- H Cl Half Equations + - H+H+ e - H Oxidation can be written as 2Cl -  Cl 2 + 2e -

82. Electroplating

83. Electroplating is when one metal is covered with a thin layer of another metal using electrolysis. Reasons to electroplate:  Protect the metal from corrosion/scratching  Improve appearance whilst saving money

84.  Plating metal used as positive electrode and is also in the solution (electrolyte)  Object to be plated is used as negative electrode  Positive electrode: plating metal loses electrons to become positive metal ions  Negative electrode: Positive metal ions gain electrons to become pure metal, plating the object + Method +

85. Negative electrode: Ni 2+ + 2e -  Ni REDUCTION Ni 2+ Electroplating copper with nickel Positive electrode: Ni - 2e -  Ni 2+ OXIDATION 2e - Half Equations + - Ni Ni 2+ Ni

86. Acids, Alkalis and Bases

87. Acids:  contain H + ions  Nitric acid (HNO 3 )  Sulfuric acid (H 2 SO 4 )  Hydrochloric acid (HCl)  Turn indicators red Alkalis:  contain OH - ions  Sodium hydroxide (NaOH)  Potassium hydroxide (KOH)  Turn indicators blue/purple

88. Alkalis are all part of a bigger group called bases. All alkalis are soluble in water but some bases are insoluble. Common bases:  Metal hydroxides  Metal oxides  Ammonia (NH 3 ) – this gives ammonium salts such as the fertiliser ammonium nitrate, NH 4 NO 3

89. The pH scale shows us the strength of an acid/alkali: pH cab be found using universal indicator: 1234567891011121314 Strong acids neutral Strong alkalis Weak acids Weak alkalis

90. Neutralisation

91. +  baseacidsalt + water Neutralisation Reactions The OH - (from base) reacts with the H + (from acid) to form the H 2 0: H + + OH -  H 2 O Examples HCl + KOH  KCl + H 2 O H 2 SO 4 + 2NaOH  Na 2 SO 4 + 2H 2 O

92. Hydrogen gives a ‘’squeaky pop’’ when it is lit Example Magnesium + Sulphuric Magnesium + Hydrogen acid sulphate +  acidmetal salt + hydrogen  Salts can also be made by reacting metals with acids

93. Naming Salts First part comes from metal in the base: e.g. copper oxide, sodium hydroxide Second part comes from the acid: Nitric acid  metal nitrate Sulphuric acid  metal sulphate Hydrochloric acid  metal chloride e.g. sodium + hydrochloric  sodium + water hydroxide acid chloride NaOH + HCl  NaCl + H 2 O

94. Obtaining the salt

95. If the salt is soluble, the solution must be evaporated to get the solid salt. +  alkaliacidsalt + water

96. If the salt made during neutralisation is insoluble it is called a precipitate. It gives a cloudy solid. Pb(NO 3 ) 2 (aq) + 2NaCl (aq)  PbCl 2 (s) + Na(NO 3 ) 2 (aq) The solid precipitate can be separated from the solution by filtration. Precipitation is used to remove waste products from water.