Chemistry

• JJ Thomson 


  In 1897 he devised the plum pudding model of the atom which had negative electrons embedded in a mass of positive charge.
• Ernest Rutherford 


  In 1911 he concluded there must be a tiny positive nucleus at the centre. This is called the Nucleus Model. The nucleus contained positively charged protons.
• Neils Bohr 


  In 1913 he refined the model so the electrons orbit around the nucleus in shells. This is known as the Bohr Model.
• James Chadwick 


  In 1932 he discovered that the nucleus also contains uncharged neutrons.
• Element 


  A substance that only contains one type of atom. Elements can't be broken down into simpler substances, and make up all materials by combining to form compounds. There are over 100 elements in the universe.
• Atom


  An atom is the smallest part of an element that can exist.

  Atoms of different elements have different atomic and mass numbers.

  All atoms are neutral so the number of positive protons must equal the number of negative electrons.
• Isotopes 


  Atoms of the same element that have the same number of protons but different number of neutrons, therefore different mass numbers.
• Relative Atomic Mass


  An average mass is an average mass that accounts for the abundance of different isotopes.
• Relative Atomic Mass Calculation 


  (mass of isotope1 x % of isotope1)+(mass of isotope2 x % of isotope2)

  ---------------------------------------------------------------

  100
• Electron Configuration 


  shell1 -> 2 electrons

  shell2 -> 8 electrons

  shell3 -> 2 electrons

  shell4 -> 0 electrons
• Compounds


  • Compounds are formed from elements in chemical reactions
  • A chemical change happens and at least one new product is formed
  • Elements are combined in fixed proportions
  • Compounds can only be separated by chemical means
• Mixtures

  • Mixtures are formed from elements or compounds combined by non-chemical means
  • The elements/compounds are unchanged
  • Elements/compounds in mixtures can be combined in any proportions
  • Mixtures can be separated by physical means
• Filtration 


  Used to separate : insoluble solid (residue) from a liquid (filtrate)

  Example : separating sand from water
• Crystallisation 


  Used to separate : soluble solid (solute) from a solvent

  Example : copper sulphate crystals from water

  The mixture is heated until the solvent starts to boil and a saturated solution is formed. As it cools, the solute becomes less soluble and some of it starts to crystallise. The remaining solvent can be left to evaporate, leaving behind solid crystals.
• Chromatography 


  Used to separate : multiple soluble substances in the same solvent

  Example : inks food dye

  A drop of the mixture is placed on chromatography paper, which is placed in a solvent. Over time the solvent soaks up the paper, and the different substances in the mixture move at different speeds depending on how strongly they interact with the paper. This causes them to separate.
• Distillation 


  Used to separate : solvent from solution

  Example : pure water from a salt solution

  The mixture is heated and the solvent boils. The vapour passes through the condenser, cools and condenses as a liquid in a collection vessel. The solute is left in the original flask.
• Half Equations 


  Only show the loss or gain of electrons of a single species during a reaction.

  Mg + 2e --> Mg2+
• Ionic Equation 


  These equations only shows species that lose or gain electrons - any that don't change are called 'spectator ions' are ignored.

  Zn + Cu2+ --> Zn2+ + Cu
• Periodic Table - Group 1 Elements 


  • Alkali metals
  • All very reactive
  • All have metallic structures
  • Reactivity increases down the group because the outer electron is further from the nucleus and easier to lose
• Periodic Table - Group 0 Elements 


  • Noble gases
  • Full outer shell of electrons so very unreactive
  • All have 8 outer electrons except Helium (2)
  • Boiling point increases down the group as the number of electron shells increases
• Periodic Table - Group 7 Elements 


  • Halogens
  • All have 7 electrons in outer shell and gain 1 more when they react
  • All reactive
  • All form diatomic molecules
  • Melting/Boiling point increases down the group
  • Reactivity decreases down the group as the outer electrons are further from the nucleus
  • React with metals to form Ionic Compounds
  • React with non-metals to form Covalent Molecules
• Metals


  Metals :

  • High melting/boiling point
  • High density
  • High malleability
  • Basic properties of oxides
  • Is conductive
  • Positive ions
• Non-Metals


  Non-Metals :

  • Low melting/boiling points
  • Low density
  • Low malleability
  • Acidic properties of oxides
  • Not conductive
  • Negative ions
• Ionic Bonding 


  • Metal + Non-Metal
  • Involve electrostatic attractions between oppositely charged ions
  • Formed when electrons are lost from one atom and transferred to another atom
  • An ion is attracted to all the oppositely charged ions in its surroundings so giant ionic lattices form
  • The formation of an ionic bond can be shown in a dot and cross diagram
• Chemical Bonds


  Involves electrostatic attraction between particles in a material. The attractive forces can be generated by the transfer or sharing of electrons between atoms. There are three main types of bonding :

  ionic, covalent, metallic.
• Metallic Bonding 


  • In a metal, atoms are arranged in giant, regular, layered structures
  • The outermost electrons of each atom are delocalised and free to move through the whole structure
  • Metallic bonding arises from electrostatic attractions between the positive nuclei of the atoms and the negative electrons
• Covalent Bonding


  • Electrons are shared between bonding atoms, not transferred
  • Bonding arises from electrostatic attractions between the positive nuclei and the bonding pair of electrons between them
  • Atoms can covalently bond to form giant lattice structures, polymers and small molecules
  • Covalent bonds can also be shown in a dot and cross diagram
• Ionic Structures

  • Giant ionic lattices are regular structures with electrostatic forces between opposite charges in all directions
  • Melting/Boiling points are high because the ionic bonds are very strong and there is lots of them
  • Solid ionic structures don't conduct electricity because there are no delocalised electrons and the ions aren't free to move
  • Dissolved/Molten ionic compounds do conduct electricity because the ions are free to move
• Metallic Structures 


  • Giant metallic lattices have ordered layers of atoms and electrostatic forces between nuclei and delocalised electrons in all directions
  • Melting/Boiling points are high because the metallic bonds are very strong
  • The layers can slide over each other making metals easy to shape
  • Delocalised electrons are free to move so metals are thermally and electrically conductive
• Diamond 


  • Giant lattice structure
  • Tetrahedral structure
  • No delocalised electrons
  • Rigid 3D structure
  • Very hard and strong
  • High melting point
  • No electrical conductivity
• Graphite


  • Giant lattice structures
  • Flat hexagonal structure
  • Layers can slide over each other
  • High melting point
  • Soft
  • Slippery
  • Electrically Conductive
• Graphene 


  • Single layer of graphite
  • Layer is only one atom thick
  • Each C atom has one delocalised electron
  • Very flexible
  • Very strong
  • Semi see-through
  • Electrically conductive
• Relative Formula Mass


  The sum of relative atomic masses of all the atoms in the formula of a compound.
• Percentage Yield 


  mass of product obtained

  ------------------------ x 100 %

  mass of product expected
• Atom Economy 


  relative formula mass of desired product

  ------------------------------------- x 100

  relative formula mass of all reactants
• Concentration


  mass

  concentration = -------

  volume
• Moles 


  1 mole = 6.02 x 10^23 particles

  The mass of one mole of a substance is equal to its relative formula mass in grams.

  mass

  moles = --------------------

  relative formula mass
• Oxidation and Reduction 


  Oxidation = Gain of oxygen

  Reduction = Loss of oxygen