All you need_to_know_about_additional_science[1]

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All you need_to_know_about_additional_science[1]

  1. 1. All you need to know about Additional Science
  2. 2. Chapters in this unit• 1. Structures and bonding• 2. Structures and properties• 3. How much?• 4. Rates of reaction• 5. Energy and reactions• 6. Electrolysis• 7. Acids, alkalis and salts
  3. 3. 1.1 Atomic structureType of sub-atomic particle Relative charge MassProton +1 1Neutron 0 1Electron -1 Negligible
  4. 4. 1.1 Atomic structure Columns = groups Group number = number of electrons in outer shellRow number = number of shells Rows = periods
  5. 5. 1.1 Atomic structureAtomic Massnumber: number:The Thenumber of number ofprotons in protonsan atom and neutrons in an atom
  6. 6. 1.2 Electronic arrangement Each shell = different energy level Shell nearest nucleus = lowest energy level Energy needed to overcome attractive forces between protons and electrons
  7. 7. 1.2 Electronic arrangement Group 1 metals (aka alkali metals) - Have 1 electron in outer most shell - Soft metals, easily cut - Reacts with water and oxygen - Reactivity increases down the group - Low melting and boiling points
  8. 8. 1.2 Electronic arrangementGroup 0/8 metals (aka noble gases)- Have 2/8 electrons in outer most shell- Very stable gases, no reaction
  9. 9. 1.2 Electronic arrangementsNo. Element Shell No. Element Shell 1 2 3 4 1 2 3 41 Hydrogen 1 11 Sodium 2 8 12 Helium 2 12 Magnesium 2 8 23 Lithium 2 1 13 Aluminium 2 8 34 Berylium 2 2 14 Silicon 2 8 45 Boron 2 3 15 Phosphorus 2 8 56 Carbon 2 4 16 Sulphur 2 8 67 Nitrogen 2 5 17 Chlorine 2 8 78 Oxygen 2 6 18 Argon 2 8 89 Fluorine 2 7 19 Potassium 2 8 8 110 Neon 2 8 20 Calcium 2 8 8 2
  10. 10. 1.3 Chemical bonding• MixtureThe combined substances do not changeEasy to separate• CompoundChemical reaction takes placeBonds form between atoms
  11. 11. 1.4 Ionic bonding (metal + non-metal)Look! Look!Group 1 element Group 7 element Very strong forces of attraction between positive and negative ions = ionic bond
  12. 12. 1.4 Ionic bonding (metal + non-metal)Ionic bonds form a giant lattice structure
  13. 13. 1.5 Covalent bonding (non-metal + non-metal)Simple molecules Giant structures
  14. 14. 1.6 Bonding in metals
  15. 15. 2.1 – 2.4 Properties Ionic Simple Giant Metallic (covalent) (covalent)Melting point intermolecular electrostatic    bonds, weak covalent covalent Strong forcesBoiling point Strong Strong forces bonds   Electrical/ Yes, when No, due No – diamond Yes, due toheat conductor molten or in to no overall Yes – delocalised solution (aq) charge graphite electrons as allows due to ions to move delocalised electrons
  16. 16. 2.3 GraphiteLayers of graphite slip off and leave a mark on paperThe free e- from each C atom can move in betweenthe layers, making graphite a good conductor ofelectricity
  17. 17. 2.4 Metal • Pure metals are made up of layers of one type of atoms • These slide easily over one another and therefore metals can be bent and shaped
  18. 18. 2.5 Nanoscience• Structures are: 1-100 nm in size or a few hundred atoms• Show different properties to same materials in bulk• Have high surface area to volume ratio
  19. 19. 2.5 Nanoscience• Titanium oxide on • Silver and socks windows Silver nanoparticles Titanium oxide in socks can prevent reacts with sunshine, the fabric from which breaks down smelling dirt
  20. 20. 3.1 Mass numbersMass number – Isotopesatomic number • Same number of protons= number of neutrons • Different number of neutronsE.g. Sodium 23 – 11 = 12
  21. 21. 3.2 Masses of atoms and molesRelative atomic masses (Ar) MolesMass of atom compared to 12C • A mole of any substance always contains samee.g. Na = 23, Cl = 35.5 number of particlesRelative formula masses (Mr) - Relative atomic mass in gramsMass of a compound found byadding Ar of each element - Relative formula mass in gramse.g. NaCl = 23 + 35.5 = 58.5
  22. 22. 3.3 Percentages and formulaePercentage mass % = mass of element total mass of compoundPercentage composition / empirical formula Al Cl Mass 9 35.5 Ar 27 35.5 Moles (9/27) = 0.33 (35.5/35.5) = 1 Simplest ratio (0.33 / 0.33) = 1 (1 / 0.33) = 3 (divide by smallest number of moles) Formula AlCl3
  23. 23. 3.4 Balancing equations H2 + O2  H2OElements Elements(Right-hand side) (Left-hand side)H= H=O= O=
  24. 24. 3.4 Reacting masses 2NaOH + Cl2  NaOCl + NaCl + H2OIf we have a solution containing 100 g of sodiumhydroxide, how much chlorine gas should we passthrough the solution to make bleach? Too much, andsome chlorine will be wasted, too little and not all of thesodium hydroxide will react.
  25. 25. 3.4 Reacting masses 2NaOH + Cl2  NaOCl + NaCl + H2O 100 g ? 2NaOH Cl2Ar / Mr 80 71Ratio (80/80) = 1 (71/80) = 0.8875 1 x 100 = 100 0.8875 x 100 = 88.75Mass 100 g 88.75 g
  26. 26. 3.5 Percentage yieldVery few chemical reactions have a yield of100% because:• Reaction is reversible• Some reactants produce unexpected products• Some products are left behind in apparatus• Reactants may not be completely pure• More than one product is produced and it maybe difficult to separate the product we want
  27. 27. 3.5 Percentage yieldPercentage yield% yield = amount of product produced (g) x 100% max. amount of product possible (g)
  28. 28. 3.5 Atom economyThe amount of the starting materials that endup as useful products is called the atomeconomy% atom economy = Mr of useful product x 100% Mr of all products
  29. 29. 3.6 Reversible reactions A+B C+D = reversible reactione.g. iodine monochloride and chlorine gas: ICl + Cl2 ICl3• increasing Cl2 increases ICl3• decreasing Cl2 decreases ICl3
  30. 30. 3.7 Haber process• Fritz Haber invented the Haber process• A way of turning nitrogen in the air into ammonia 450oC 200 atmN2 + 3H2 2NH3
  31. 31. 4.2 Collision theoryCollision theory Rate of reactionChemical reactions increases if:only occur when • temperaturereacting particles increasescollide with each other • concentration orwith sufficient energy. pressure increases The minimum amount • surface areaof energy is called the increasesactivation energy • catalyst used
  32. 32. 4.2 Surface areaWhy?The inside of a large piece of solid is not incontact with the solution it is reacting with, soit cannot reactHow?Chop up solid reactant intosmaller pieces or crush intoa powder
  33. 33. 4.3 TemperatureWhy?At lower temperatures, particles will collide:a) less oftenb) with less energyHow?Put more energy into reactionIncreasing the temperatureby 10oC will double the rateof reaction
  34. 34. 4.4 ConcentrationWhy?Concentration is a measure of how manyparticles are in a solution. Units = mol/dm3The lower the concentration, the fewerreacting particles, the fewer successfulcollisionsHow?Add more reactant to thesame volume of solution
  35. 35. 4.4 PressureWhy?Pressure is used to describe particles in gasesThe lower the pressure, the fewer successfulcollisionsHow?Decrease the volume orIncrease the temperature
  36. 36. 4.5 CatalystsWhy?Expensive to increase temperature or pressureDo not get used up in reaction and can bereusedHow?Catalysts are made from transition metals, e.g.iron, nickel, platinumProvide surface area for reacting particles tocome together and lower activation energy
  37. 37. 5.1 Energy changesExothermic reaction, Endothermic reaction,e.g. respiration e.g. photosynthesis• Energy ‘exits’ • Energy ‘enters’reaction – heats reaction – coolssurroundings surroundings• Thermometer • Thermometerreadings rises readings fall
  38. 38. 5.2 Energy and reversible reactions Exothermic reaction Hydrated Anhydrouscopper sulphate copper sulphate + water Endothermic reaction
  39. 39. 5.3 Haber process (again!) Exothermic  temperature,  temperature, reaction  products  products N2 + 3H2 2NH3 Endothermic  temperature,  temperature, reaction  products  products
  40. 40. 5.3 Haber process (again!) Smaller vol. of  pressure,  pressure, gas produced  products  products N2 + 3H2 2NH3 Larger vol. of  pressure,  pressure, gas produced  products  products
  41. 41. 5.3 Haber process (again!)Temperature:- Forward reaction is exothermic, so low temperature is preferred- But this makes reaction slow- Compromise by using 450OC N2 + 3H2 2NH3 Catalyst:Pressure: - Iron- The higher the better - Speeds up both- High pressure is dangerous! sides of reaction- Compromise by using 200-350 atm
  42. 42. 6.1 Electrolysis Electrolysis: splitting up using electricity Ionic substance - molten (l) - dissolved (aq)Non-metal ionMetal ion
  43. 43. 6.2 Changes at the electrodesSolutionsWater contains the ions: H+ and O2-The less reactive element will be given off at electrode Oxidation is loss Reduction is gain OIL RIG Molten (PbBr) 2Br-  Br2 + 2e- Pb2+ + 2e-  Pb Solution (KBr) 2Br-  Br2 + 2e- 2H+ + 2e-  H2
  44. 44. 6.3 Electrolysing brine At anode 2Cl- (aq)  Cl2 (g) + 2e- At cathode 2H+ (aq) + 2e-  H2 (g) In solution Na+ and OH-
  45. 45. 6.4 Purifying copper At anode 2H2O (l)  4H+ (aq) + O2 (g) + 2e- At cathode Cu2+ (aq) + 2e-  Cu (s)
  46. 46. 7.1 Acids and alkalisAcids = H+ ionsAlkalis = OH- ionsAlkalis = soluble bases
  47. 47. 7.2 + 7.3 Salts Acid Formula Salt ExampleHydrochloric HCl Chloride Sodium chlorideSulphuric H2SO4 Sulphate Copper sulphateNitric HNO3 Nitrate Potassium nitrate
  48. 48. 7.2 + 7.3 Salts – metals, bases and alkalisMetals:Metal(s) + acid(aq)  salt(aq) + hydrogen(g)Bases: Acid(aq) + base(aq)  salt(aq) + water(l)Alkalis: Acid(aq) + alkali(aq)  salt(aq) + water(l)Ionic equation (neutralisation): H+ + OH-  H2O
  49. 49. 7.3 Salts – solutionsSolutions:solution(aq) + solution(aq)  precipitate(s) + solution(aq)Solid precipitate is filtered off and dried

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