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Chemistry Paper 1, It is necessary to calibrate a pH meter so as to give…

- - - - Halogens
        
        Oxidising power decreases down the group
      - Halides
        
        Reducing power increases down the group
  - - - Slow N/A
        
        Al2O3
        
        Amphoteric
        
        7
    - - Slow N/A
        
        SiO2
        
        Acid
        
        7
    - - Vigorous brill. white flame
        
        MgO
        
        Base
        
        8-9
    - - Spont. combust brill. white flame
        
        P4O10
        
        Acid
        
        2-4
    - - Vigorous yellow flame
        
        Na2O
        
        Base
        
        12-14
        
        Basic due to O2- : O2- + H2O ➔ 2OH-
    - - Steady purple flame
        
        SO2/SO3
        
        Acid
        
        2-4/1-3
  - - - OH- = Mg least soluble
      - SO42 = Mg most soluable
  - - - All elements within a period have the same number of electron shells (if you don't worry about s and p sub-shells)
      - All the elements within a group have the same number of electrons in their outer shell-so they have similar properties.
      - The group number tells you the number of electrons in the outer shell.
      - The periodic table can be split into an s block, d block, p block and f block like this: Doing this shows you which sub shells all the electrons go into.
- - - - ∆H, whilst important, is not sufficient to explain feasible change.
      - The concept of increasing disorder (entropy change, ∆S).
      - ∆S accounts for the above deficiency, illustrated by physical changes and chemical changes.
        The balance between entropy and enthalpy determines the feasibility of a reaction given by the relationship:
        
        ∆G = ∆H – T∆S
        
        For a reaction to be feasible, value of ∆G must be zero or negative
    - - Lattice enthalpy can be defined as either enthalpy of lattice dissociation or enthalpy of lattice formation.
      - Born-Haber cycles are used to calculate lattice enthalpies using the following data:
        
        Enthalpy of formation
        
        Ionisation energy
        
        Bond enthalpy
        
        Enthalpy of atomisation
        
        Electron affinity
  - - - Water is slightly disssociated
      - Kw derived from the equilibrium constant for this dissociation
      - Kw=[H+] [OH-]
      - The value of Kw varies with temperature
    - - weak acids and bases dissociate only slightly in aqueous solution
      - Ka is the dissociation constant for a weak acid
      - pKa=-log Ka
    - - pH=-log [H+]
      - The concentration of hydrogen ions in aqueous solution covers a very wide range. Therefore, a logarithmic scale, the pH, is used as a measure of hydrogen ion concentration
    - - A buffer solution maintains an approximately constant pH, despite dilution or addition of small amounts of acid or base.
      - Acidic buffer solutions contain a weak acid and the salt of that weak acid.
      - Basic buffer solutions contain a weak base and the salt of that weak base.
      - Applications of buffer solutions
    - - An acid is a proton donor
      - A base is a proton acceptor
      - Acid-base equilibria involve the transfer of protons
  - - - A loss of electrons=Oxidation
      - A gain in electrons=Reduction
      - Reduction and Oxidation happen at the same time = term redox reaction
      - Oxidising agent accepts electrons and gets reduced its self
      - Reducing agent donates electrons and gets oxidised its self
    - - In compound ions, the overall oxidation state is just the ion charge
        
        SO4)2- overall oxidation state =-2,oxidation state of O=-2 (total=4x-2=-8), so oxidation state of S=-2-2(-8)=+6
      - The sum of the states for a neutral compound is 0
        
        Fe2O3 overall oxidation state=0 Oxidation state of O=-2 (total=3x-2=-6) so oxidation state of Fe=(0-(-6))÷2=+3
      - The oxidation state of a simple monatomic ion, like Na+ is the same as its charge
      - Combined oxygen is nearly always -2, except in peroxides, where it's -1 (and in fluorides OF2, where its +2, and O2F2, where it's +1 and O2 where it's O)
        
        In H2O, oxidation state of O=-2, but in H2O2, oxidation state of O=-1
      - Elements just bonded to identical atoms, like O2 and H2, also have an oxidation state of 0
      - Combined hydrogen is +1, except in metal hydrides where it is -1 (and H2 where it's 0)
        
        In HF, oxidation state of H=+1, but in NaH, oxidation state of H+-1
      - Uncombined elements, eg He and Ar, have an oxidation state of 0
    - - Sometimes, oxidation states aren't clear from the formula of a compound. If you see Roman numerals in a chemical name, it's an oxidation number. E.g. iron has oxidation state +2 in iron(II) sulfate (FeSO4), but it has oxidation state +3 in iron(III) sulfate (Fe2(SO4)3)
    - - You can write Half-equations and combine them into Redox Equations
      - Ionic half-equations show oxidation or reduction.
      - You can show the electrons that are being lost or gained in a half equation. For example, this is the half-equation for the oxidation of sodium: Na➔Na+ + e-
      - You can combine half-equations for different oxidising or reducing agents together to make full equations for redox reactions.
  - - - Mass number (A)
      - Atomic (proton) number (Z)
      - How do you determine the number of fundamental particles in atoms and ions using mass number, atomic number and charge ?
      - Explain the existence of isotopes ?
      - The principles of a TOF mass spectrometer, limited to ionisation, acceleration to give all ions constant kinetic energy, ion drift, ion detection, data analysis
      - The mass spectrometer gives accurate information about relative isotopic mass and also about the relative abundance of isotopes.
      - Mass spectrometry can be used to identify elements
      - Mass spectrometry can be used to determine relative molecular mass
    - - Define first ionisation energy?
        
        The energy needed to remove 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions.
      - Write equation for first successive ionisation energy
        
        O➔O+ + e-
      - Explain how first and successive ionisation energies in period 3 (Na-Ar) and in Group 2 (Be-Ba) give evidence for electron configuration in sub-shells and in shells?
      - Factors affecting ionisation energy
        
        Nuclear charge
        
        The more protons there are in the nucleus, the more positively charged the nucleus is and the stronger the attraction for the electrons.
        
        Distance from nucleus
        
        Attraction falls off very rapidly with distance. An electron close to the nucleus will be much more strongly attached than one further away.
        
        Shielding
        
        As the number of electrons between the outer electrons and the nucleus increases, the outer electrons feel less attraction to the nucleus. This lessening of the pull of the nucleus thanks to the inner electron shells is called shielding.
      - Electronic structure
    - - Timeline
        
        1913
        
        Bohr-Nuclear model (with shells)
        
        Electrons have set positions in atoms called shells
        
        1911
        
        Rutherford-Nuclear model
        
        Region of positive charge in centre of atom with region of negative charge around the nucleus. Discovered by gold foil experiment
        
        1932
        
        James Chadwick-Discovers the neutron
        
        1897
        
        J.J Thompson-"Atoms aren't solid spheres"
        
        1803
        
        John Dalton-"Atoms are solid sphere"
      - Particles
        
        Proton
        
        relative charge=+1
        
        relative mass=1
        
        Electron
        
        relative charge=-1
        
        relative mass=1/2000
        
        Neutron
        
        relative charge=0
        
        relative mass=1
      - Atom consists of a nucleus containing protons and neutrons surrounded by electrons
- - - - Cations
        
        Group 2 NH4+
      - Anions
        
        Group 7 (halide ions),OH-,CO32-,SO42-
    - - You can mix chlorine with water it undergoes disproportionation
        
        You end up with a mixture of chloride ions and chlorate(I) ions
      - 1) In the UK our drinking water is treated to make sure it is safe. Chlorine is an important part of water treatment:
        
        It kills disease-causing microorganisms
        
        Some chlorine persists in the water and prevents reinfection further down the supply
        
        It prevents the growth of algae, eliminating bad tastes and smells, and removes discolouration caused by organic compounds.
      - 2)However, there are risks from using chlorine to treat water
        
        Chlorine gas is very harmful if it's breathed in - it irritates the respiratory system. Liquid chlorine on the skin or eyes causes severe chemical burns. Accidents involving chlorine could be really serious, or fatal.
        
        Water contains a variety of organic compounds, e.g. from the decomposition of plants. Chlorine reacts with these compounds to form chlorinated hydrocarbons, e.g. chloromethane (CH3Cl), and many of these chlorinated hydrocarbons are carcinogenic. However this increased cancer risk is small compared to the risks from untreated water - a cholera epidemic, say, could kill thousands of people.
      - Chlorate(I) ions kill bacteria. So adding chlorine (or a compound containing chlorate (I) ions) to water can make it safe to drink or swim in. On the downside, chlorine is toxic :cry:
      - We have to weigh up these risks and benefits when making decisions about whether we should add chemicals to drinking water supplies.
  - - - Phenolphthalein-Colourless in acid and pink in base
      - Methyl orange-Red in acid and yellow in base
    - - 1) Rinse the burette (ethanoic acid), and fill with same solution label burette
      - 2) Use burette to transfer 20cm3 to a clean 100cm3 beaker
      - 3) Rinse a second burette with other solution (NaOH) then fill this burette
      - 4) Rinse pH probe with distilled/deionised water and use rod to stir solution and record pH reading in a suitable table
      - 5) Add in 2cm3 portions until 18cm3 and 0.2cm3 until 22cm3 and 2cm3 until 40cm3
      - 6) Rinse the pH probe with distilled/deionised water after readings
      - Analysing data
        
        Use calibrating curve to adjust data collected from acid base reaction
        
        Then plot graph of adjusted pH
        
        Now you can interpret there was strong or weak solutions