Foundations in Chemistry

Foundations in Chemistry

1. Atoms Ions & Compounds

   • Atomic Structure and isotopes
     • Definitions
       • Isotopes: atoms of the same element with different numbers of neutrons and different masses
       • Relative isotopic mass: mass of an element compared with 1/12th mass of carbon 12
       • relative atomic mass: weighted mean mass of an element compared with 1/12th mass of carbon 12
   • Relative mass
     • Definitions
       • Mass Defects: the mass lost by the strong nuclear forces holding protons and neutrons together
     • 1 u is approximately the mass of 1 proton or neutron
       • u is the unit for atomic mass unit
       • carbon-12 is exactly 12 u
     • mass to charge ratio M/Z = relative mass of an ion/relative charge on an ion
     • weighted mean mass takes into account 1) The percentage abundance of each isotope 2) relative isotonic mass of each isotope
   • The Mole
     • 1 mole is 6.02*10^23 of something
       • avogadros constant
     • Number of particles = mass (g) / molar mass (g mol^-1)

2. Amounts of Substance

   • Amount of Substance
     • Gas
       • Volume Moles x 24
       • 1 mole of gas has a volume of 24dm^3
     • Solid
       • Mass Moles x Mr
   • Determination of formulae
     • Definitions
       • Empirical formula: the simplest smallest whole number ratio of atoms of each element in a compound
     • Mr = Ar+Ar+Ar
     • Hydrated salts
       • 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
       • Assumptions have to be made
         • 1) All the water is lost - to make sure heat repeatedly until weight no longer changes, showing all the water has been removed
         • 2) No further decomposition - many salts decompose further when heated, which would be difficult to notice if theres no colour change
   • Moles and Volumes
     • In a solution
       • Moles = concentration x volume
       • May need to convert from cm to dm^3
       • A standard solution is a solution of a known concentration
         • Made by dissolving an exact mass of solute a solvent and making it to an exact volume
     • A molar gas volume is the volume per mole of gas at a stated temperature and pressure
       • RTP
         • 101KPa
         • 20ºC
         • Molar gas volume is always 24 dm^3 at RTP
       • Volume = moles x 24
       • Ideal Gas Equation
         • 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
   • Reacting Quantities
     • Percentage yield
       • Maximum possible is theoretical yield
       • 1) reaction may not have been completed
       • 2) Side reactions may have taken place
       • 3) Purification of the product may cause a loss in product
     • Limiting Reagents
       • Reagents that is not in excess
       • Whatever has the most moles as it will be used up first
     • Atom Economies
       • How well atoms have been utilised
         • If high then lots of desired products and sustainable
         • Atom economy = sum of molar mass of desired products / sum of molar masses of all products

3. Acids and Redox

   • Acids Bases & Neutralisation
     • A strong acid releases all its H+ ions &completely dissociates in aqueous solution
       • e.g. HCl
     • A weak acid only releases a proportion of it's H+ atoms as it only partially dissociates
       • e.g Ethanoic acid
         • most organic acids are weak
     • Bases neutralises an acid to form a salt
       • e.g. metal oxides/hydroxides/carbonates & ammonia
       • An Alkali is a base that dissolves in water realising hydroxide ions into the solution
       • Acid+Base=Salt+Water
       • Only Carbonates produce anything other than water and salt (CO2)
   • Acid-Base titrations
     • Can be used for
       • Identifying chemicals
       • finding concentrations
       • finding purity of a substance
         • important for quality control
           • e.g medicine food ect.
     • Volumetric flasks are used to create standard solutions
     • Titration Calculations
       • You will know V1 and C1 and V2
       • 1) Work out the mol of the solution that you know C1 and V1
       • 2) Use ratios to find the mol of the other substance
       • 3) Use the mol and V2 to work out C2
   • Redox
     • Definitions
       • Oxidation Number: a measure of the number of electrons that an atoms uses to bond with atoms of another element.
     • Oxidation Numbers
       • Pure elements are always 0
       • Sign is placed before the number
       • O = -2
       • H = +1
       • F = -1
       • Will match charge
         • e.g. Mg2+ = +2
       • Special cases
         • H in hydrides = -1
         • O in peroxides = -1
         • O bonded to F = +2
       • Sum of Oxidation Numbers = Total charge
       • Group 1 = +1
       • Group 2 = +2
       • Al = +3
     • Roman numerals shows the oxidation state of the element without a sign
       • Which will be the same as the charge
       • NO2- is nitrite
         • NO3- is nitrate
           • SO4^2- is sulfate
     • OIL RIG
       • of electrons
       • gaining oxygen is oxidation
       • Reduction is decrease to oxidation number
         • Oxidation is increase in oxidation number
     • Some Dilute Acids undergo redox reactions with metals to produce salts and hydrogen gas

4. Electrons and Bonding

   • Electron strucure
     • Shells are energy levels within the atom. The energy increase as the shell number (principle quantum number0 increases
     • 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.
       • S-orbitals can hold 1 or 2 electrons, the greater the wheel number, the greater its s-orbital radius, spherical shape.
       • There are 3 P-orbitals, each holding 2 electrons, have a figure of 8 shape.
     • Sub-shells are within a shell, and are the same type of orbital.
     • Orbitals fill in order of increasing energy
       • 4s fills before 3d
     • electrons are negatively charged so repel each other, so have opposite spins in an orbital, shown by up and down arrows
       • Orbitals with the same energy are occupied singly first
     • 4s fills first, but is also emptied first, as once 3d is filled its energy falls below 4s
   • Ionic bonding & Structure
     • Ionic bonding is the electrostatic attraction between positive and negative ions
       • Cations include metal ions & ammonium ions
       • Anions include non-metal ions & polyatomic ions
       • Dot and cross diagrams using square brackets
       • Each ion attracts oppositley in all directions, so an ionic compound has a giant ionic lattice structure
         • 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.
         • 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
         • 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
   • Covalent Bonding
     • Covalent bonding is the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms
       • A covalent bond is the overlap of atomic orbitals, each one containing one electron to give a shared pair of electrons
         • The shared pair of electrons is attracted to each of the nuclei
         • Both have the same structure in the outer shell as the nearest noble gas.
     • Occurs with non-metal elements/compounds and polyatomic ions
     • The attraction is localised acting only between the shared pair of electrons, resulting in a small unit called a molecule
     • Number of bonds
       • Carbon - 4
       • Nitrogen - 3
       • Oxygen - 2
       • Hydrogen - 1
     • Multiple bonds - when two atoms share more than 1 pair of electrons
       • Double bonds
     • Dative or Coordinate bonds are were one atom pzrovidesthe shared pair of electrons which was originally a lone pair, e.g. NH4+

5. Shapes of Molecules and intermolecular forces

   • Shapes of Molecules
     • Electrons has negative charges so repel one another
       • determines shape of molecule
       • lone pair is slightly closer to the central atom & occupies more space so repels more strongly than a bonded pair
     • Angles
       • 2 bonded pairs, linear, 180º
       • 3 bonded pairs, trigonal planar, 120º
       • 4 bonded pairs, tetrahedral, 109.5º
       • 6 bonded pairs, octahedral, 90º
       • 3 bonded arts, 1 lone pair, pyramidal, 107º
       • 2 bonded pairs, 2 lone pairs, non-linear, 104.5º
   • Electronegativity & Polarity
     • Definition
       • Electronegativity is the measure of attraction of a bonded atom for the pair of electrons in a covalent bond
     • FONCl has the highest Pauling electronegativity value, With having the highest of 4.0
     • Bond Type Electronegativity Difference
     • Covalent 0
     • Polar Covalent 0 to 1.8
     • Ionic greater than 1.8