AS

• How do ID-ID bonds occur?

  Electrons are in continuous random motion, and may be distributed unevenly, creating an instantaneous dipole. This induces a dipole on a neighbouring molecule, creating an induced dipole. There is an electrostatic attraction between the two dipoles.
• What factors affect the strength of ID-ID bonds?

  • Number of electrons in the molecule (e.g Mr)
  • Distance between molecules (e.g increased branching decreases the surface area in contact)
  • Chain length (e.g longer chains increases the number of ID-ID bonds)
• Requirements for hydrogen bonding

  • A delta + H atom in one molecule bonded to an electronegative atom (O, N or F)
  • A small electronegative atom in the other molecule with a lone pair
• Group 2 hydroxides increase in solubility down the group
• Group 2 carbonates decrease in solubility down the group
• Lithium, sodium, potassium, and ammonium salts are soluble
• Nitrates are soluble
• Most halides are soluble, except:

  • silver halides
  • copper iodide (white precipitate)
  • lead chloride and bromide (white precipitate)
  • lead iodide (yellow precipitate)
• Most sulfates are soluble except:

  • barium
  • calcium
  • lead (ALL WHITE PRECIPITATES)
• Most hydroxides are INSOLUBLE except:

  • Lithium
  • Sodium
  • Potassium
  • Strontium
  • Calcium
  • Barium
  • Ammonium
• Most carbonates are insoluble except:

  • Sodium, Lithium, Potassium and Ammonium
• Colours of insoluble carbonates:

  • Cu = blue-green
  • Ag = yellow
  • Most others are white
• Electromagnetic spectrum in order of increasing frequency and decreasing wavelength

  Infra-red, visible light, Ultraviolet
• How are absorption or emission spectra produced?

  • Electrons can absorb energy from their surroundings causing them to be excited to higher energy levels.
  • They only absorb certain frequencies corresponding to differences between energy levels, meaning the radiation passing through has certain frequencies missing
  • Electrons emit energy when dropping down to lower energy levels, and the energy levels are discrete so these spectra will be unique to different atoms.
• Enthalpy change of formation Hess’s cycle

  Arrows point out of elements (at the bottom) to products and reactants (at the top)
• Enthalpy changes of combustion Hess’s cycle

  The arrows point towards combustion products at the bottom
• Cracking can take place at lower temperatures by passing hydrocarbon vapour over a heated solid catalyst
• NOx are produced in car engines and can react with unburnt hydrocarbons in the prescence of sunlight to form ground level ozone, which is a component of photochemical smog- which can irritate the eyes and respiratory system.
• Producing alkanes from alkenes using H2

  • A nickel catalyst and 150’C
  • Platinum catalyst and RTP
• • Electrophiles are electron pair acceptors.
  • They are attracted to areas of electron density
  • Examples: positively charged ions, polar molecules
• Mechanism for electrophilic addition (e.g of alkene and bromine):

  • Br—Br is polarised and heterolytic fission occurs
  • The closer positive Br gives up the bonding electrons to the other Br
  • The electron pair from the double bond moves to the positive Br forming a bond to one carbon, and a carbocation intermediate on the other carbon
  • The Br- then bonds to the other carbon atom
• How can alcohols be produced using alkenes?

  • Hydration of ethene using an acid catalyst: Add cold concentrated H2SO4, reacts in an electrophilic addition reaction. Then add cold water and warm the product and it is hydrolysed to form an alcohol
  • Hydrating ethene with steam at 300’C and high pressure, using a solid H3PO4 catalyst
• Disadvantages of biofuels:

  • Car engines have to be modified
  • Land used to grow crops for fuel can’t be used to grow food
  • Still produce CO2 when burnt
• Electrolysis of an aqueous solution products at the cathode:

  • Metal is less reactive than hydrogen = metal will be formed (e.g silver or copper)
  • Metal more reactive (e.g group 1/2/Al) = hydrogen formed (2H+ + 2e- —> H2
• Electrolysis of an aqueous solution products at the anode:

  • No halide = oxygen formed from hydroxide ions in the water ( 4OH- —> O2 + 2H2O + 4e-)
  • Halide = halogen produced
• If bromide is displaced (e.g by Cl2) and Br2 is formed, the reaction mixture will turn orange. If you add hexane there will be an orange/red layer above.

  If iodine is displaced then the reaction mixture will turn brown and a pink/violet layer will form when hexane is added
• Hydrogen halides can be made using an ionic halide and concentrated sulfuric acid

  (e.g Cl- + H3PO4 —> HCl + H2PO4-)
• • H2SO4 is an oxidising agent and I2 and Br2 are strong enough reducing agents to reduce H2SO4
  • You can make HCl using H2SO4 (NaCl + H2SO4 —> HCl + HSO4-) but you can’t make HBr or HCl.
  • Bromide ions are oxidised to bromine gas and the sulfuric acid is reduced to sulfur dioxide ( ionic equation: H2SO4 + 2H+ + 2Br- —>Br2 + SO2 + 2H2O)
  • Iodine reduces it to H2S (H2SO4 + 8H+ + 8I- —> 4I2 + H2S + 4H2O)
• The thermal stability of hydrogen halides decreases down the group because the bond enthalpy gets weaker as the halogen atoms get bigger
• Hydrogen halides reactions with sulfuric acid:

  • HBr reduces H2SO4 to SO2: 2HBr + H2SO4 —> Br + SO2 + 2H2O
  • HI reduces it to H2S: 8HI + H2SO4 —> 4I2 + H2S + 4H2O
• • Kc >> 1 = equilibrium lies far to the right
  • Kc << 1 = equilibrium lies far to the left
• The boiling points of the haloalkanes increases down the group because the atomic radius increases and the number of electron shells increases, leading to stronger ID-ID forces
• A nucleophile is an electron pair donor which attacks areas which are electron deficient, for example OH-, NH3 and H2O.
• Haloalkanes can go under nucleophilic substitution (e.g to form an alcohol or amine) because a nucleophile can attack the delta + carbon and replace the delta - atom or group (the halogen)
• Mechanism for nucleophilic substitution

  • Haloalkanes + warm NaOH under reflux (R-X + NaOH —> ROH + NaX)
  • The C-Br bond is polar, and the delta +C attracts a lone pair of electrons from the OH- ion
  • The C-Br bond breaks heterolytically and a new bond forms between the C and OH-, leaving the Br
  • This can occur with water too (forming an intermediate with a O+ so one of the O-H bonds breaks)
• Haloalkanes react with ammonia by electrophilic substitution to form amines. In the second step another ammonia removes a hydrogen from the N+H3 group to leave an amine.

  
• • Halogen radicals react with alkanes to form haloalkanes: CH4 + Cl2 –uv—> CH3Cl +HCl
  • Initiation: photodissociation of Cl2, homolytic fission to form 2 Cl radicals
  • Cl radical attacks a methane molecule forming a hydrogen halide (HCl) and methyl radical (CH3)
  • Methyl radical can attack another Cl2 molecule to form CH3Cl and another Cl radical
  • If 2 free radicals join they have a termination reaction (e.g CH3 + Cl)

  A) Photodissociation

  B) Homolytic fission
• • Ozone is formed in the stratosphere when UV radiation from the sun hits oxygen molecules (O2 + hv —> O + O, O2 + O —> O3)
  • UV radiation can also reverse the formation of ozone so the equilibrium set up is O2 + O <—> O3
  • The ozone layer is important because when it breaks down it absorbs harmful high energy UV radiation which can cause skin cancer/the DNA in cells/sun burns
• Ozone can occur in the troposphere (lower part of atmospher) due to the effect of sunlight on mixtures of NO2 and hydrocarbons.

  Ground level ozone mixes with particulates to create photochemical smog which can cause respiratory problems.
• • CFCs have C-Cl bonds which can be broken down by high energy UV in the stratosphere to form Cl radicals which act as catalysts in the breakdown of ozone.
  • CCl3 + hv —> CCl2F + Cl
  • They react with ozone to form a ClO intermediate and O2
  • Cl + O3 —> O2 + ClO
  • ClO + O —> O2 + Cl
  • So the overall reaction is O3 + O —> O2
  • In the troposphere only some CFCs are broken down because most high energy UV has been absorbed by the ozone layer so less ozone is broken down in the troposphere.