Foundations in Chemistry

Atoms Ions & Compounds
1. Atomic Structure and isotopes
  1. Definitions
    1. Isotopes: atoms of the same element with different numbers of neutrons and different masses
    2. Relative isotopic mass: mass of an element compared with 1/12th mass of carbon 12
    3. relative atomic mass: weighted mean mass of an element compared with 1/12th mass of carbon 12
2. Relative mass
  1. Definitions
    1. Mass Defects: the mass lost by the strong nuclear forces holding protons and neutrons together
  2. 1 u is approximately the mass of 1 proton or neutron
    1. u is the unit for atomic mass unit
    2. carbon-12 is exactly 12 u
  3. mass to charge ratio M/Z = relative mass of an ion/relative charge on an ion
  4. weighted mean mass takes into account 1) The percentage abundance of each isotope 2) relative isotonic mass of each isotope
3. The Mole
  1. 1 mole is 6.02*10^23 of something
    1. avogadros constant
  2. Number of particles = mass (g) / molar mass (g mol^-1)
Amounts of Substance
1. Amount of Substance
  1. Gas
    1. Volume Moles x 24
    2. 1 mole of gas has a volume of 24dm^3
  2. Solid
    1. Mass Moles x Mr
2. Determination of formulae
  1. Definitions
    1. Empirical formula: the simplest smallest whole number ratio of atoms of each element in a compound
  2. Mr = Ar+Ar+Ar
  3. Hydrated salts
    1. blue colour from water, when water is removed so is the colour so only a white powder remain that is anhydrous, but still slightly blue, as difficult to remove all the water
    2. Assumptions have to be made
      1. 1) All the water is lost - to make sure heat repeatedly until weight no longer changes, showing all the water has been removed
      2. 2) No further decomposition - many salts decompose further when heated, which would be difficult to notice if theres no colour change
3. Moles and Volumes
  1. In a solution
    1. Moles = concentration x volume
    2. May need to convert from cm to dm^3
    3. A standard solution is a solution of a known concentration
      1. Made by dissolving an exact mass of solute a solvent and making it to an exact volume
  2. A molar gas volume is the volume per mole of gas at a stated temperature and pressure
    1. RTP
      1. 101KPa
      2. 20ºC
      3. Molar gas volume is always 24 dm^3 at RTP
    2. Volume = moles x 24
    3. Ideal Gas Equation
      1. pV = nRT
        P is Pressure in pascals
        V is volume in m cubed
        n is moles of gas
        R is ideal gas constant = 8,31 J mol^-1 K^-1
        T is Temperature in K
4. Reacting Quantities
  1. Percentage yield
    1. Maximum possible is theoretical yield
    2. 1) reaction may not have been completed
    3. 2) Side reactions may have taken place
    4. 3) Purification of the product may cause a loss in product
  2. Limiting Reagents
    1. Reagents that is not in excess
    2. Whatever has the most moles as it will be used up first
  3. Atom Economies
    1. How well atoms have been utilised
      1. If high then lots of desired products and sustainable
      2. Atom economy = sum of molar mass of desired products / sum of molar masses of all products
Acids and Redox
1. Acids Bases & Neutralisation
  1. A strong acid releases all its H+ ions &completely dissociates in aqueous solution
    1. e.g. HCl
  2. A weak acid only releases a proportion of it's H+ atoms as it only partially dissociates
    1. e.g Ethanoic acid
      1. most organic acids are weak
  3. Bases neutralises an acid to form a salt
    1. e.g. metal oxides/hydroxides/carbonates & ammonia
    2. An Alkali is a base that dissolves in water realising hydroxide ions into the solution
    3. Acid+Base=Salt+Water
    4. Only Carbonates produce anything other than water and salt (CO2)
2. Acid-Base titrations
  1. Can be used for
    1. Identifying chemicals
    2. finding concentrations
    3. finding purity of a substance
      1. important for quality control
        e.g medicine food ect.
  2. Volumetric flasks are used to create standard solutions
  3. Titration Calculations
    1. You will know V1 and C1 and V2
    2. 1) Work out the mol of the solution that you know C1 and V1
    3. 2) Use ratios to find the mol of the other substance
    4. 3) Use the mol and V2 to work out C2
3. Redox
  1. Definitions
    1. Oxidation Number: a measure of the number of electrons that an atoms uses to bond with atoms of another element.
  2. Oxidation Numbers
    1. Pure elements are always 0
    2. Sign is placed before the number
    3. O = -2
    4. H = +1
    5. F = -1
    6. Will match charge
      1. e.g. Mg2+ = +2
    7. Special cases
      1. H in hydrides = -1
      2. O in peroxides = -1
      3. O bonded to F = +2
    8. Sum of Oxidation Numbers = Total charge
    9. Group 1 = +1
    10. Group 2 = +2
    11. Al = +3
  3. Roman numerals shows the oxidation state of the element without a sign
    1. Which will be the same as the charge
    2. NO2- is nitrite
      1. NO3- is nitrate
        SO4^2- is sulfate
  4. OIL RIG
    1. of electrons
    2. gaining oxygen is oxidation
    3. Reduction is decrease to oxidation number
      1. Oxidation is increase in oxidation number
  5. Some Dilute Acids undergo redox reactions with metals to produce salts and hydrogen gas
Electrons and Bonding
1. Electron strucure
  1. Shells are energy levels within the atom. The energy increase as the shell number (principle quantum number0 increases
  2. An Atomic Orbital is a region around the nucleus that can hold two electron of opposite spins. they are visualised as areas of probability for finding an electron.
    1. S-orbitals can hold 1 or 2 electrons, the greater the wheel number, the greater its s-orbital radius, spherical shape.
    2. There are 3 P-orbitals, each holding 2 electrons, have a figure of 8 shape.
  3. Sub-shells are within a shell, and are the same type of orbital.
  4. Orbitals fill in order of increasing energy
    1. 4s fills before 3d
  5. electrons are negatively charged so repel each other, so have opposite spins in an orbital, shown by up and down arrows
    1. Orbitals with the same energy are occupied singly first
  6. 4s fills first, but is also emptied first, as once 3d is filled its energy falls below 4s
2. Ionic bonding & Structure
  1. Ionic bonding is the electrostatic attraction between positive and negative ions
    1. Cations include metal ions & ammonium ions
    2. Anions include non-metal ions & polyatomic ions
    3. Dot and cross diagrams using square brackets
    4. Each ion attracts oppositley in all directions, so an ionic compound has a giant ionic lattice structure
      1. Most are solids at room temperature.
        There is insufficient energy to overcome the strong electrostatic attraction between the oppositely charged ions. High angry levels are required so they have a high melting point
        Higher melting point for greater attractions.
      2. Many ionic compounds dissolve in polar solvents such as water.
        Polar molecules break down the lattice and surround each ion in solution
        In compounds with larger charges, the attraction may be too strong to overcome so won't be very soluble
      3. As a solid ionic compounds don't conduct electricity as the ons are fixed in position so there are no mobile charge carriers
        As a liquid or when dissolved in an aqueous solution the ionic lattice has broken down, so the ions can act as mobile charge carriers
3. Covalent Bonding
  1. Covalent bonding is the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms
    1. A covalent bond is the overlap of atomic orbitals, each one containing one electron to give a shared pair of electrons
      1. The shared pair of electrons is attracted to each of the nuclei
      2. Both have the same structure in the outer shell as the nearest noble gas.
  2. Occurs with non-metal elements/compounds and polyatomic ions
  3. The attraction is localised acting only between the shared pair of electrons, resulting in a small unit called a molecule
  4. Number of bonds
    1. Carbon - 4
    2. Nitrogen - 3
    3. Oxygen - 2
    4. Hydrogen - 1
  5. Multiple bonds - when two atoms share more than 1 pair of electrons
    1. Double bonds
  6. Dative or Coordinate bonds are were one atom pzrovidesthe shared pair of electrons which was originally a lone pair, e.g. NH4+
Shapes of Molecules and intermolecular forces
1. Shapes of Molecules
  1. Electrons has negative charges so repel one another
    1. determines shape of molecule
    2. lone pair is slightly closer to the central atom & occupies more space so repels more strongly than a bonded pair
  2. Angles
    1. 2 bonded pairs, linear, 180º
    2. 3 bonded pairs, trigonal planar, 120º
    3. 4 bonded pairs, tetrahedral, 109.5º
    4. 6 bonded pairs, octahedral, 90º
    5. 3 bonded arts, 1 lone pair, pyramidal, 107º
    6. 2 bonded pairs, 2 lone pairs, non-linear, 104.5º
2. Electronegativity & Polarity
  1. Definition
    1. Electronegativity is the measure of attraction of a bonded atom for the pair of electrons in a covalent bond
  2. FONCl has the highest Pauling electronegativity value, With having the highest of 4.0
  3. Bond Type Electronegativity Difference
  4. Covalent 0
  5. Polar Covalent 0 to 1.8
  6. Ionic greater than 1.8

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