Chemistry 2

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  • Chemistry 2
    • Analytical Techniques
      • Infra-red spectroscopy
        • The peaks are where the radiation is absorbed by the bonds.
      • Mass Spectrometry
        • 2 peaks with heights in a 3:1 ratio = Chlorine.
        • 2 peak at the same height =Bromine
    • Group 2
      • Mg, Ca, Sr, Ba
      • Size of atoms
        • Atoms increase in size as the group descends.
      • Melting Points
        • Have quite high melting points because of their metallic structures
          • Positive ions surrounded by a sea of delocalised electrons
        • Mg is lower because it has a different crystal structure
      • First Ionisation energies
        • As the group descends, this decreases because there is more shielding so it is easier to remove the 1st electron.
      • Reactions with water
        • Ba + H2O --> Ba(OH)2 + H2
          • General equation
      • Solubility of hydroxides
        • As the group descends the solubility increases.
      • Solubility of sulphates
        • As the group descends the solubility increases.
      • Testing for sulphate ions
        • Add dilute HCl
          • Add barium chloride
            • White precipitate (Barium Sulphate) will form in sulphate ions present.
          • Removes carbonate ions so they dont give precipitate.
          • Sulphuric acid cannot be used (would give precipitate)
      • Uses of group 2
        • Mg(OH)2 - An alkali which you can drink to neutralise stomach acid.
        • Ca(OH)2 - Lime. Can be spread on fields and will neutralise acidic soil.
        • BaSO4 - If drank it will block x-rays. So can be used to take scans of digestive system.
          • Ba2+ - It is lethal, but cannot be absorbed because it is so insoluble
    • Oxidation and Reduction
      • Oxidation is LOSS of electrons
        • Oxidising agents - Electron ACCEPTERS
      • Reduction is GAIN of electrons.
        • Reducing agents - Electron DONORS
      • Oxidation states
        • Group 1 = +1
        • Group 2 = +2
        • Al = +3
        • Ions = Charge on ion
        • Oxygen = -2
          • If in peroxide = -1
        • Hydrogen = +1
          • Unless metal hydride = -1
        • Group 7 = -1
        • Elements by themselves = 0
    • Polymerisation of alkenes
      • Formed by the addition of monomers to the end of a growing chain.
      • Temp = 200 degrees
      • Pressure = 2000 atmospheres
    • Stereoisomers
      • Definition: Isomers which have the same structural formula but atoms arranged differently in space.
      • If the molecules are on the same (ZAME) side then a 'Z' is put infront of the name.
      • If molecules are on the opposite side then an 'E' is put infront of the name.
      • The bigger the carbon chain, the more important it is.
    • Alcohols
      • Fermentation
        • This uses yeast to convert glucose into ethanol and carbon dioxide.
        • At low temps the reaction is very slow and at high temps the enzymes denature, so the process is carried out at 35 degrees.
        • Product = Biofuel.
        • C6H12O6 ----> 2C2H5OH + 2CO2
        • Expensive on manpower but cheap equipment, impure, batch, renewable.
      • Direct Hydration
        • The industrial process, the direct hydration of ethene using steam and a phosphoric acid catalyst at 300 degrees and 6.5x10^3 kPa.
        • C2H4(g) + H2O(g) ---> C2H5OH(g)
        • Cheap on manpower but expensive equipment, pure, continuous, non-renewable.
      • Ethanol is considered to be CARBON NEUTRAL in that the amount of carbon dioxide released when burnt is equal to that which the plants absorb during photosynthesis.
      • Primary Alcohol: One R group joined to the carbon with the OH on it.
      • Secondary Alcohol: 2 R groups joined to the carbon with the OH on it.
      • Tertiary Alcohol: 2 R groups joined to the carbon with the OH on it.
      • Oxidation of Alchols
        • Products: Aldehyde, Ketone, Carboxylic Acid.
        • Oxidising agents: Kr2Cr2O7, Na2Cr2O7.
        • Primary alcohols are oxidised into aldehydes and carboxylic acids.
        • Secondary alcohols are oxidised into ketones.
        • During the course of these reactions we see a colour change from ORANGE to GREEN.
        • Test for aldehyde and ketone: Add TOLLEN'S REAGENT and the aldehyde will produce a SILVER MIRROR.
          • No observation with ketones.
    • Kinematics
      • Maxwell-Boltzmann distribution
        • Effect of Catalysts: The Ea line will move to the left. No effect on the shape of the curve.
          • Increase the amount of successful collisions.
        • Effect of Temperature: As the temp. increases the curve will stretch to the right and shorten.
          • Increase the amount of successful collisions.
      • Increase Rate of Reaction.
        • TEMPERATURE - increases the speed of the particles and therefore the number of successful collisions.
        • CONCENTRATION - increases the number of particles and therefore the number of successful collisions.
        • PRESSURE - same as concentration.
        • SURFACE AREA - exposes more sites available for collision, so more successful collisions.
        • CATALYSTS - lowers the activation energy therefore previously unsuccessful collisions are now successful.
    • Extraction of Metals.
      • Using CARBON.
        • 1) Good reducing agent. 2) Cheap. 3) Plentiful. 4) Good in the extraction of iron.
        • Blast furnace
          • Reaction 1: Coke reacts with hot air in an exothermic reaction. This produces the heat needed for the reduction of iron(III) oxide.
            • C(s) + O2(g) --> CO2(g)
          • Reaction 2: CO2 reacts at high temps. with more coke to form CO.
            • CO2(g) + C(s) --> 2CO(g)
          • Reaction 3: The CO then reduces most of the iron(III) oxide at around 1200 degrees.
            • Fe2O3(s) + 3CO(g) --> 2Fe(l) + 3CO2(g)
          • Reaction 4: In the hotter part of the furnace the iron oxide can also be reduced directly using the coke.
        • The extraction of Manganese
          • Produces IMPURE manganese.
          • Heat manganese(IV) oxide with carbon.
          • MnO2(s) + C(s) --> Mn(l) + CO2(g)
        • The extraction of Copper.
          • Heat copper(II) oxide with carbon.
          • Produces impure copper.
          • 2CuO(s) +C(s) --> 2Cu(l) +CO2(g).
      • Metal Sulphides.
        • They are problematic because they can produce SO2 which is toxic and causes acid rain.
        • SO2 can be converted to H2SO4 .
        • The sulphide is roasted with oxygen.
          • 2ZnS + 3O2 --> 2ZnO +2SO2
      • Tungsten.
        • Reducing with carbon; the product is tungsten carbide which is useless and brittle.
        • We use H2 as a reducing agent; very explosive so great care has to be taken.
        • WO3 + 3H2 --> W(s) + 3H2O
      • Titanium.
        • TiO2 + 2Cl2 + C --> TiCl4 + CO2
          • TiCl4 + 4Na --> 4NaCl + Ti
          • TiCl4 + 2Mg --> 2MgCl2 +Ti
        • Batch Process.
        • Problems: Expensive, inert.
      • Aluminium.
        • Does not react with carbon so we use electrolysis.
        • AlO3 - Bauxite/Cryolyte reduces melting point.
        • Al3+ + 3e- --> Al(l)
          • 2O(2-) --> O2 + 4e-
            • O2 + C --> CO2
    • Equilibria
      • Concentration
        • Whichever concentration we INCREASE the equilibrium will move to the opposite side.
      • Pressure
        • If we INCREASE the pressure; equilibrium will move to the side with the LEAST MOLES.
      • Le Chatellier's Principle states: If a system at equilibrium is disturbed then the equilibrium position will move in the direction that undoes the disturbance.
      • Dynamic equilibrium = The rate of the forward reaction is equal to the rate of the backward reaction in a closed system.
      • Catalysts
        • They DO NOT affect the equilibrium position. Just increase rate of reaction.
      • Temperature
        • If we INCREASE the temp....
          • Endothermic...
            • Equilibrium will move to the right.
          • Exothermic...
            • Equilibrium will move to the left.
    • Enthalpy
      • Exothermic = - delta H.
      • Endothermic = + delta H.
      • Standard enthalpy change: occur at 100kPa and stated temp, usually 298K.
      • Standard enthalpy of formation: Enthalpy change when 1 mole of a compound is formed from its elements under standard conditions.
        • Enthalpy of formation = Product - Reactants.
      • Hess's Law: Enthalpy change of a reaction depends only on the initial and final states of the reaction and is independant of the route taken.
      • Standard enthalpy of combustion: Enthalpy change when 1 mole of a compound is completely burned in oxygen under standard conditions.
        • Enthalpy of combustion = Reactants - Products.
    • Group 7
      • The Halogens
        • F, Cl, Br, I, At.
        • F - Pale green gas/ Cl - pale green gas/ Br - brown liquid/ I - Black solid.
        • Atomic Radius: As the group descends the atom radius increases.
        • Electronegativity: As the group descends there is an increase in shielding and nuclear charge. So Fluorine is the most electronegative because it has the least shielding.
        • Boiling Points: We have strong covalent bonds holding the molecules together but to boil we need to overcome the van der waals. The more electrons, the bigger the van der waals. Therefore boiling points increase as you go down the group.
        • When dissolved in water:
          • Fluorine - Very pale green.
          • Chlorine - Very pale green.
          • Bromine - Yellow
          • Iodine - Brown with black precipitate.
      • Halides
        • eg. Fluoride
          • Always colourless when dissolved in water.
        • Testing for halides: Add nitric acid to get rid of carbide atoms, then add silver nitrate solution.
          • Fluoride: Colourless --> Colourless
          • Chloride: Colourless --> White precipitate --> dissolves in dilute ammonia.
            • Ag(+)(aq) + CL(-)(aq) --> AgCl(s)
          • Bromide: Colourless --> Cream precipitate --> Dissolves in conc. ammonia.
          • Iodide: Colourless --> Yellow precipitate.
        • Reducing Ability of Halides.
          • Reactions with conc. H2SO4.
          • Fluoride/Chloride do not reduce conc. H2SO4 so only the hydrogen halide made - Steamy fumes.
          • Bromide reduces a little so SO2(choking smell), steamy fumes. Brown fumes from bromide getting oxidised.
          • Iodide reduces a lot so SO2(choking smell), S(yellow solid), H2S(rotten eggs smell). Iodide gets oxidised so black solid/purple fumes.
      • Reaction with Halogens/Halides
        • F2 - best oxidising agent.
        • I2 - Worst oxidising agent.
        • The smaller the atom the greater the energy to break the bond.
          • F2 an exception because the two lone pairs repel each other which weakens the bond, so best oxidising agent.
        • Gain an electron - energy is released, exothermic, similar amount for each halogen.
        • Energy is released when H from water molecules are attracted to the ion. Attraction is stronger is ion is smaller.
  • Equilibria
    • Concentration
      • Whichever concentration we INCREASE the equilibrium will move to the opposite side.
    • Pressure
      • If we INCREASE the pressure; equilibrium will move to the side with the LEAST MOLES.
    • Le Chatellier's Principle states: If a system at equilibrium is disturbed then the equilibrium position will move in the direction that undoes the disturbance.
    • Dynamic equilibrium = The rate of the forward reaction is equal to the rate of the backward reaction in a closed system.
    • Catalysts
      • They DO NOT affect the equilibrium position. Just increase rate of reaction.
    • Temperature
      • If we INCREASE the temp....
        • Endothermic...
          • Equilibrium will move to the right.
        • Exothermic...
          • Equilibrium will move to the left.

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