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AQA Chemistry C2 Revision


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AQA GCSE Chemistry C2 Revision notes

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AQA Chemistry C2 Revision

  1. 1. 26/04/14 AQA Chemistry 2AQA Chemistry 2 A slideshow that covers the entire AQA 2006 Syllabus Chemistry 2 Module W Richards
  2. 2. 26/04/14 The structure of the atomThe structure of the atom ELECTRON – negative, mass nearly nothing PROTON – positive, same mass as neutron (“1”) NEUTRON – neutral, same mass as proton (“1”) The Ancient Greeks used to believe that everything was made up of very small particles. I did some experiments in 1808 that proved this and called these particles ATOMS: Dalton
  3. 3. 26/04/14 Mass and atomic numberMass and atomic number Particle Relative Mass Relative Charge Proton 1 +1 Neutron 1 0 Electron Very small -1 MASS NUMBER = number of protons + number of neutrons SYMBOL PROTON NUMBER = number of protons (obviously)
  4. 4. 26/04/14 Mass and atomic numberMass and atomic number How many protons, neutrons and electrons?
  5. 5. 26/04/14 IsotopesIsotopes An isotope is an atom with a different number of neutrons: Each isotope has 8 protons – if it didn’t then it just wouldn’t be oxygen any more. Notice that the mass number is different. How many neutrons does each isotope have?
  6. 6. 26/04/14 Electron structureElectron structure Consider an atom of Potassium: Potassium has 19 electrons. These electrons occupy specific energy levels “shells”… Nucleus The inner shell has __ electrons The next shell has __ electrons The next shell has __ electrons The next shell has the remaining __ electron Electron structure = 2,8,8,1
  7. 7. 26/04/14 Periodic Table IntroductionPeriodic Table Introduction
  8. 8. 26/04/14 Mendeleev Periodic tablePeriodic table The periodic table arranges all the elements in groups according to their properties. Horizontal rows are called PERIODS Vertical columns are called GROUPS
  9. 9. 26/04/14 H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Fe Ni Cu Zn Br Kr Ag I Xe Pt Au Hg The Periodic TableThe Periodic Table Fact 1: Elements in the same group have the same number of electrons in the outer shell (this corresponds to their group number) E.g. all group 1 metals have __ electron in their outer shell These elements have __ electrons in their outer shell These elements have __ electrons in their outer shells
  10. 10. 26/04/14 H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Fe Ni Cu Zn Br Kr Ag I Xe Pt Au Hg The Periodic TableThe Periodic Table Fact 2: As you move down through the periods an extra electron shell is added: E.g. Lithium has 3 electron in the configuration 2,1 Potassium has 19 electrons in the configuration __,__,__,__ Sodium has 11 electrons in the configuration 2,8,1
  11. 11. 26/04/14 H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Fe Ni Cu Zn Br Kr Ag I Xe Pt Au Hg The Periodic TableThe Periodic Table Fact 3: Most of the elements are metals: These elements are metals This line divides metals from non- metals These elements are non-metals
  12. 12. 26/04/14 H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Fe Ni Cu Zn Br Kr Ag I Xe Pt Au Hg The Periodic TableThe Periodic Table Fact 4: (Most important) All of the elements in the same group have similar PROPERTIES. This is how I thought of the periodic table in the first place. This is called PERIODICITY. E.g. consider the group 1 metals. They all: 1) Are soft 2) Can be easily cut with a knife 3) React with water
  13. 13. 26/04/14 CompoundsCompounds Compounds are formed when two or more elements are chemically combined. Some examples: Glucose Methane Sodium chloride (salt)
  14. 14. 26/04/14 Some simple compounds…Some simple compounds… Methane, CH4 Water, H2O Carbon dioxide, CO2 Ethyne, C2H2 Sulphuric acid, H2SO4 Key Hydrogen Oxygen Carbon Sulphur
  15. 15. 26/04/14 Balancing equationsBalancing equations Consider the following reaction: Na O H H H H Na O H Sodium + water sodium hydroxide + hydrogen + + This equation doesn’t balance – there are 2 hydrogen atoms on the left hand side (the “reactants” and 3 on the right hand side (the “products”)
  16. 16. 26/04/14 Balancing equationsBalancing equations We need to balance the equation: Na O H H H H Na O H Sodium + water sodium hydroxide + hydrogen + + Na O H H Na O H Now the equation is balanced, and we can write it as: 2Na(s) + 2H2O(l) 2NaOH(aq) + H2(g)
  17. 17. 26/04/14 Some examplesSome examples Mg + O2 Zn + HCl Fe + Cl2 NaOH + HCl CH4 + O2 Ca + H2O NaOH + H2SO4 CH3OH + O2 MgO ZnCl2 + H2 FeCl3 NaCl + H2O CO2 + H2O Ca(OH)2 + H2 Na2SO4 + H2O CO2 + H2O 2 2 2 3 2 2 2 2 3 2 2 2 2 2 4
  18. 18. 26/04/14 BondingBonding Hi. My name’s Johnny Chlorine. I’m in Group 7, so I have 7 electrons in my outer shell I’d quite like to have a full outer shell. To do this I need to GAIN an electron. Who can help me? Cl Cl
  19. 19. 26/04/14 BondingBonding Here comes one of my friends, Harry Hydrogen Hey Johnny. I’ve only got one electron but it’s really close to my nucleus so I don’t want to lose it. Fancy sharing? Cl H Cl H Now we’re both really stable. We’ve formed a covalent bond.
  20. 20. 26/04/14 BondingBonding Here comes another friend, Sophie Sodium Hey Johnny. I’m in Group 1 so I have one electron in my outer shell. Unlike Harry, this electron is far away from the nucleus so I’m quite happy to get rid of it. Do you want it? Cl Now we’ve both got full outer shells and we’ve both gained a charge. We’ve formed an IONIC bond. Na Okay Cl Na +-
  21. 21. 26/04/14 Covalent bondingCovalent bonding Consider an atom of hydrogen: Notice that hydrogen has just __ electron in its outer shell. A full (inner) shell would have __ electrons, so two hydrogen atoms get together and “_____” their electrons: Now they both have a ____ outer shell and are more _____. The formula for this molecule is H2. When two or more atoms bond by sharing electrons we call it ____________ BONDING. This type of bonding normally occurs between _______ atoms. It causes the atoms in a molecule to be held together very strongly but there are ____ forces between individual molecules. This is why covalently-bonded molecules have low melting and boiling points (i.e. they are usually ____ or ______). Words – gas, covalent, non-metal, 1, 2, liquid, share, full, weak, stable
  22. 22. 26/04/14 Dot and Cross DiagramsDot and Cross Diagrams H OH Water, H2O:
  23. 23. 26/04/14 Dot and Cross DiagramsDot and Cross Diagrams Oxygen, O2: O O
  24. 24. 26/04/14 Dot and cross diagramsDot and cross diagrams Water, H2O: Oxygen, O2: OH H O O H H O O O Step 1: Draw the atoms with their outer shell: Step 2: Put the atoms together and check they all have a full outer shell:
  25. 25. 26/04/14 Dot and cross diagramsDot and cross diagrams Nitrogen, N2: Carbon dioxide, CO2:Ammonia NH3: Methane CH4: H HN H HH H H C N N O OC
  26. 26. 26/04/14 Other ways of drawing covalent bondsOther ways of drawing covalent bonds Consider ammonia (NH3): H HN H H HN H H HN H
  27. 27. 26/04/14 IonsIons An ion is formed when an atom gains or loses electrons and becomes charged: If we “take away” the electron we’re left with just a positive charge: This is called an ion (in this case, a positive hydrogen ion) + - + The electron is negatively charged The proton is positively charged +
  28. 28. 26/04/14 Ionic bondingIonic bonding Na Na + This is where a metal bonds with a non-metal (usually). Instead of sharing the electrons one of the atoms “_____” one or more electrons to the other. For example, consider sodium and chlorine: Sodium has 1 electron on its outer shell and chlorine has 7, so if sodium gives its electron to chlorine they both have a ___ outer shell and are ______. A _______ charged sodium ion A _________ charged chloride ion As opposed to covalent bonds, ionic bonds form strong forces of attraction between different ions due to their opposite ______, causing GIANT IONIC STRUCTURES to form (e.g sodium chloride) with ______ melting and boiling points: Cl Cl -
  29. 29. 26/04/14 Some examplesSome examples Mg Magnesium chloride: MgCl2 Cl Cl + Mg 2+ Cl - Cl - Calcium oxide: CaO OCa + Ca 2+ O 2-
  30. 30. 26/04/14 Giant structures (Giant structures (““latticeslattices””)) + + + + +++ + + 1. Diamond – a giant covalent structure with a very ____ melting point due to ______ bonds between carbon atoms 2. Graphite – carbon atoms arranged in a layered structure, with free _______ in between each layer enabling carbon to conduct _________ 3. Sodium chloride – a giant ionic lattice with _____ melting and boiling points due to ______ forces of attraction. Can conduct electricity when _______. 4. Metals – the __________ in metals are free to move around (“delocalised”), holding the _____ together and enabling it to conduct _________
  31. 31. 26/04/14 A closer look at metalsA closer look at metals + + + + +++ + + Metals are defined as elements that readily lose electrons to form positive ions. There are a number of ways of drawing them: + - + - + - +- + - + - + - + - + + + + + + + + Delocalised electrons
  32. 32. 26/04/14 NanoscienceNanoscience Nanoscience is a new branch of science that refers to structures built from a few hundred atoms and are 1- 100nm big. They show different properties to the same materials in bulk. They also have a large surface area to volume ratio and their properties could lead to new developments in computers, building materials etc. Definition: Task: research nanoscience and find two current and/or future applications of this science.
  33. 33. 26/04/14 Group 1 – The alkali metalsGroup 1 – The alkali metals Li Na K Rb Cs Fr
  34. 34. 26/04/14 Group 1 – The alkali metalsGroup 1 – The alkali metals 1) These metals all have ___ electron in their outer shell. Some facts… 2) Reactivity increases as you go _______ the group. This is because the electrons are further away from the _______ every time a _____ is added, so they are given up more easily. 3) They all react with water to form an alkali (hence their name) and __________, e.g: Words – down, one, shell, hydrogen, nucleus, decreases Potassium + water potassium hydroxide + hydrogen 2K(s) + 2H2O(l) 2KOH(aq) + H2(g) 2) Density increases as you go down the group, while melting point ________
  35. 35. 26/04/14 Group 0 – The Noble gasesGroup 0 – The Noble gases He Ne Ar Kr Xe Rn
  36. 36. 26/04/14 Group 0 – The Noble gasesGroup 0 – The Noble gases Some facts… 1) All of the noble gases have a full outer shell, so they are very ______ 2) They all have _____ melting and boiling points 3) They exist as single atoms rather then _________ molecules 4) Helium is ________ then air and is used in balloons and airships (as well as for talking in a silly voice) 5) Argon is used in light bulbs (because it is so unreactive) and argon , krypton and ____ are used in fancy lights Words – neon, stable, low, diatomic, lighter
  37. 37. 26/04/14 Group 7 – The halogensGroup 7 – The halogens F Cl Br I At
  38. 38. 26/04/14 Group 7 – The HalogensGroup 7 – The Halogens Some facts… 1) Reactivity DECREASES as you go down the group Decreasing reactivity (This is because the electrons are further away from the nucleus and so any extra electrons aren’t attracted as much). 2) They exist as diatomic molecules (so that they both have a full outer shell): Cl Cl 3) Because of this fluorine and chlorine are liquid at room temperature and bromine is a gas
  39. 39. 26/04/14 The halogens – some reactionsThe halogens – some reactions 1) Halogen + metal: Na + Cl - Na Cl+ 2) Halogen + non-metal: H Cl+ Cl H Halogen + metal ionic salt Halogen + non-metal covalent molecule
  40. 40. 26/04/14 Atomic massAtomic mass SYMBOL PROTON NUMBER = number of protons (obviously) RELATIVE ATOMIC MASS, Ar (“Mass number”) = number of protons + number of neutrons
  41. 41. 26/04/14 Relative formula mass, MRelative formula mass, Mrr The relative formula mass of a compound is the relative atomic masses of all the elements in the compound added together. E.g. water H2O: Therefore Mr for water = 16 + (2x1) = 18 Work out Mr for the following compounds: 1) HCl 2) NaOH 3) MgCl2 4) H2SO4 H=1, Cl=35 so Mr = 36 Na=23, O=16, H=1 so Mr = 40 Mg=24, Cl=35 so Mr = 24+(2x35) = 94 H=1, S=32, O=16 so Mr = (2x1)+32+(4x16) = 98 Relative atomic mass of O = 16 Relative atomic mass of H = 1
  42. 42. 26/04/14 AA ““MoleMole”” Definition: A mole of a substance is the relative formula mass of that substance in grams For example, 12g of carbon would be 1 mole of carbon... ...and 44g of carbon dioxide (CO2) would be 1 mole etc...
  43. 43. 26/04/14 Calculating percentage massCalculating percentage mass If you can work out Mr then this bit is easy… Calculate the percentage mass of magnesium in magnesium oxide, MgO: Ar for magnesium = 24 Ar for oxygen = 16 Mr for magnesium oxide = 24 + 16 = 40 Therefore percentage mass = 24/40 x 100% = 60% Percentage mass (%) = Mass of element Ar Relative formula mass Mr x100% Calculate the percentage mass of the following: 1) Hydrogen in hydrochloric acid, HCl 2) Potassium in potassium chloride, KCl 3) Calcium in calcium chloride, CaCl2 4) Oxygen in water, H2O
  44. 44. 26/04/14 Recap questionsRecap questions Work out the relative formula mass of: 1) Carbon dioxide CO2 2) Calcium oxide CaO 3) Methane CH4 Work out the percentage mass of: 1) Carbon in carbon dioxide CO2 2) Calcium in calcium oxide CaO 3) Hydrogen in methane CH4
  45. 45. 26/04/14 Calculating the mass of a productCalculating the mass of a product E.g. what mass of magnesium oxide is produced when 60g of magnesium is burned in air? Step 1: READ the equation: 2Mg + O2 2MgO IGNORE the oxygen in step 2 – the question doesn’t ask for it Step 3: LEARN and APPLY the following 3 points: 1) 48g of Mg makes 80g of MgO 2) 1g of Mg makes 80/48 = 1.66g of MgO 3) 60g of Mg makes 1.66 x 60 = 100g of MgO Step 2: WORK OUT the relative formula masses (Mr): 2Mg = 2 x 24 = 48 2MgO = 2 x (24+16) = 80
  46. 46. 26/04/14 Work out Mr: 2H2O = 2 x ((2x1)+16) = 36 2H2 = 2x2 = 4 1. 36g of water produces 4g of hydrogen 2. So 1g of water produces 4/36 = 0.11g of hydrogen 3. 6g of water will produce (4/36) x 6 = 0.66g of hydrogen Mr: 2Ca = 2x40 = 80 2CaO = 2 x (40+16) = 112 80g produces 112g so 10g produces (112/80) x 10 = 14g of CaO Mr: 2Al2O3 = 2x((2x27)+(3x16)) = 204 4Al = 4x27 = 108 204g produces 108g so 100g produces (108/204) x 100 = 52.9g of Al O 1) When water is electrolysed it breaks down into hydrogen and oxygen: 2H2O 2H2 + O2 What mass of hydrogen is produced by the electrolysis of 6g of water? 3) What mass of aluminium is produced from 100g of aluminium oxide? 2Al2O3 4Al + 3O2 2) What mass of calcium oxide is produced when 10g of calcium burns? 2Ca + O2 2CaO
  47. 47. 26/04/14 Another methodAnother method Try using this equation: Mass of product IN GRAMMES Mass of reactant IN GRAMMES Mr of product Mr of reactant Q. When water is electrolysed it breaks down into hydrogen and oxygen: 2H2O 2H2 + O2 What mass of hydrogen is produced by the electrolysis of 6g of water? Mass of product IN GRAMMES 6g 4 36 So mass of product = (4/36) x 6g = 0.66g of hydrogen
  48. 48. 26/04/14 Problems with this techniqueProblems with this technique Calculating the amount of a product may not always give you a reliable answer... 1) The reaction may not have completely _______ 2) The reaction may have been _______ 3) Some of the product may have been ____ 4) Some of the reactants may have produced other _______ The amount of product that is made is called the “____”. This number can be compared to the maximum theoretical amount as a percentage, called the “percentage yield”. Words – lost, yield, finished, reversible, products
  49. 49. 26/04/14 Atom EconomyAtom Economy Percentage atom economy = Relative formula mass of useful product Total masses of products Calculate the atom economies of the following: 1) Converting ethanol into ethene (ethene is the useful bit): C2H5OH C2H4 + H20 2) Making zinc chloride from zinc and hydrochloric acid: Zn + 2HCl ZnCl2 + H2
  50. 50. 26/04/14 Numbers of molesNumbers of moles Consider two liquids: Now consider two gases: 20cm3 of 0.1mol/dm3 of hydrochloric acid 20cm3 of 0.1mol/dm3 of sodium hydroxide These two beakers contain the same number of moles 20cm3 of helium at room temperature and pressure 20cm3 of argon at room temperature and pressure These two gases contain the same number of moles
  51. 51. 26/04/14 Endothermic and exothermic reactionsEndothermic and exothermic reactions 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 Energy Energy
  52. 52. 26/04/14 Example reactionsExample reactions Reaction Temp. after mixing/O C Exothermic or endothermic? Sodium hydroxide + dilute hydrochloric acid Sodium hydrogencarbonate + citric acid Copper sulphate + magnesium powder Sulphuric acid + magnesium ribbon
  53. 53. 26/04/14 Reversible ReactionsReversible Reactions Some chemical reactions are reversible. In other words, they can go in either direction: A + B C + D NH4Cl NH3 + HCl e.g. Ammonium chloride Ammonia + hydrogen chloride If a reaction is EXOTHERMIC in one direction what must it be in the opposite direction? For example, consider copper sulphate: Hydrated copper sulphate (blue) Anhydrous copper sulphate (white) + Heat + Water CuSO4 + H2OCuSO4.5H2O
  54. 54. 26/04/14 Reversible ReactionsReversible Reactions When a reversible reaction occurs in a CLOSED SYSTEM (i.e. no reactants are added or taken away) an EQUILIBRIUM is achieved – in other words, the reaction goes at the same rate in both directions: A + B C + D Endothermic reactions Increased temperature: Decreased temperature: A + B C + D A + B C + D More products Less products Exothermic reactions Increased temperature: Decreased temperature: A + B C + D Less products More products A + B C + D
  55. 55. 26/04/14 Making AmmoniaMaking Ammonia Nitrogen + hydrogen Ammonia N2 + 3H2 2NH3 •High pressure •450O C •Iron catalyst Recycled H2 and N2 Nitrogen Hydrogen Mixture of NH3, H2 and N2. This is cooled causing NH3 to liquefy. Fritz Haber, 1868-1934 Guten Tag. My name is Fritz Haber and I won the Nobel Prize for chemistry. I am going to tell you how to use a reversible reaction to produce ammonia, a very important chemical. This is called the Haber Process. To produce ammonia from nitrogen and hydrogen you have to use three conditions:
  56. 56. 26/04/14 Haber Process: The economicsHaber Process: The economics A while ago we looked at reversible reactions: A + B C + D Endothermic, increased temperature A + B C + D Exothermic, increase temperature ExothermicEndothermic 1) If temperature was DECREASED the amount of ammonia formed would __________... 2) However, if temperature was INCREASED the rate of reaction in both directions would ________ causing the ammonia to form faster 3) If pressure was INCREASED the amount of ammonia formed would INCREASE because there are less molecules on the right hand side of the equation Nitrogen + hydrogen Ammonia N2 + 3H2 2NH3
  57. 57. 26/04/14 Haber Process SummaryHaber Process Summary •200 atm pressure •450O C •Iron catalyst Recycled H2 and N2 Nitrogen Hydrogen Mixture of NH3, H2 and N2. This is cooled causing NH3 to liquefy. To compromise all of these factors, these conditions are used: A low temperature increases the yield of ammonia but is too slow A high temperature improves the rate of reaction but decreases the yield too much A high pressure increases the yield of ammonia but costs a lot of money
  58. 58. 26/04/14 Rates of ReactionRates of Reaction Hi. I’m Mike Marble. I’m about to have some acid poured onto me. Let’s see what happens… Here comes an acid particle… It missed! Here comes another one. Look at how slow it’s going… No effect! It didn’t have enough energy! Oh no! Here comes another one and it’s got more energy…
  59. 59. 26/04/14 Rates of ReactionRates of Reaction Chemical reactions occur when different atoms or molecules _____ with enough energy (the “________ Energy): Basically, the more collisions we get the _______ the reaction goes. The rate at which the reaction happens depends on four things: 1) The _______ of the reactants, 2) Their concentration 3) Their surface area 4) Whether or not a _______ is used Words – activation, quicker, catalyst, temperature, collide
  60. 60. 26/04/14 CatalystsCatalysts Task Research and find out about two uses of catalysts in industry, including: 1) Why they are used 2) The disadvantages of each catalyst
  61. 61. 26/04/14 Catalyst SummaryCatalyst Summary Catalysts are used to ____ __ a reaction to increase the rate at which a product is made or to make a process ________. They are not normally ___ __ in a reaction and they are reaction-specific (i.e. different reactions need _________ catalysts). Words – different, speed up, used up, cheaper
  62. 62. 26/04/14 Rate of reaction graph v1Rate of reaction graph v1 Time taken for reaction to complete Temperature/ concentration Reaction takes a long time here Reaction is quicker here
  63. 63. 26/04/14 Rate of reaction graph v2Rate of reaction graph v2 Amount of product formed/ amount of reactant used up Time Slower reaction Fast rate of reaction here Slower rate of reaction here due to reactants being used up Rate of reaction = amount of product formed/reactant used up
  64. 64. 26/04/14 ElectrolysisElectrolysis + + + + - - - - Positive electrode Cu2+ Cu2+ Cu2+ Negative electrode Cl- Cl- Cl- Solution containing copper and chloride ions
  65. 65. 26/04/14 ElectrolysisElectrolysis Electrolysis is used to separate a metal from its compound. = chloride ion = copper ion When we electrolysed copper chloride the _____ chloride ions moved to the ______ electrode and the ______ copper ions moved to the ______ electrode – OPPOSITES ATTRACT!!!
  66. 66. 26/04/14 Electrolysis equationsElectrolysis equations We need to be able to write “half equations” to show what happens during electrolysis (e.g. for copper chloride): 2 2 2 At the negative electrode the positive ions GAIN electrons to become neutral copper ATOMS. The half equation is: Cu2+ + e- Cu At the positive electrode the negative ions LOSE electrons to become neutral chlorine MOLECULES. The half equation is: Cl- - e- Cl2
  67. 67. 26/04/14 Purifying CopperPurifying Copper + + + + - - - - Solution containing copper ions Impure copper Cu2+ Cu2+ Cu2+ Pure copper At the positive electrode: Cu(s) Cu2+ (aq) + 2e- At the negative electrode: Cu2+ (aq) + 2e- Cu(s)
  68. 68. 26/04/14 Electrolysis of brineElectrolysis of brine Positive electrode Negative electrode Sodium chloride (brine) NaCl(aq) Sodium hydroxide (NaOH(aq)) Sodium chloride (salt) is made of an alkali metal and a halogen. When it’s dissolved we call the solution “brine”, and we can electrolyse it to produce 3 things… Chlorine gas (Cl2) Hydrogen gas (H2)
  69. 69. 26/04/14 Universal Indicator and the pH scaleUniversal Indicator and the pH scale Strong acid Strong alkali Neutral 1 2 3 4 5 6 7 8 9 10 11 12 13 14 Universal Indicator is a mixture of liquids that will produce a range of colours to show how strong the acid or alkali is: Stomach acid Lemon juice Water Soap Oven cleanerBaking powder An acid contains hydrogen ions, H+ An alkali contains hydroxide ions, OH-
  70. 70. 26/04/14 Neutralisation reactionsNeutralisation reactions When acids and alkalis react together they will NEUTRALISE each other: OHNa Sodium hydroxide ClH Hydrochloric acid The sodium “replaces” the hydrogen from HCl ClNa Sodium chloride H2O Water General equation: H+ (aq) + OH- (aq) H2O(l)
  71. 71. 26/04/14 Making saltsMaking salts Whenever an acid and alkali neutralise each other we are left with a salt, like a chloride or a sulphate. Complete the following table: Hydrochloric acid Sulphuric acid Nitric acid Sodium hydroxide Sodium chloride + water Potassium hydroxide Potassium sulphate + water Calcium hydroxide Calcium nitrate + water
  72. 72. 26/04/14 Making SaltsMaking Salts Soluble salts can be made from acids by reacting them with: 1) Metals, e.g. Zn + 2HCl ZnCl2 + H2 2) Insoluble bases, e.g. CuO + 2HCl CuCl2 + H20 3) Alkalis (alkali = a “soluble base”), e.g. NaOH + HCl NaCl + H20 Salts can be made from these solutions by crystallizing them.
  73. 73. 26/04/14 Ammonium saltsAmmonium salts Guten tag again. When ammonia dissolves in water it produces an alkaline solution: NH3 + H20 NH4OH This solution can be used to make fertilisers. Very useful!
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