1. Atoms
   1. Basic atomic properties
      1. Simple background
         1. The sub-atomic particles
         2. nucleus
            1. Topic
            2. Isotope
         3. Electron
            1. Topic
      2. Atomic Orbitals
         1. Heisenberg Uncertainty Principle
         2. 1s2, 2p6, 3d10, 4f14 orbitals
            1. Topic
            2. Topic
            3. Topic
            4. Topic
            5. Topic
         3. Hund's rule
      3. Electron structures
         1. Topic
         2. Shortcut
         3. Ca
         4. Chromium/copper: s moves to d to make single orbitals
         5. Ion structures
            1. REMEMBER
      4. Ionisation Energy
         1. Topic
         2. Factors
            1. more protons = more IE
            2. More distance from nucleus = less IE
            3. More shielding/screening = less IE
            4. paired electrons = less IE
         3. A few elements
            1. Hydrogen
            2. Helium
            3. Lithium
         4. General trends
            1. Topic
            2. Exceptions
            3. group 2-3 (Be - B)
            4. group 5-6 (N-O)
            5. Topic
            6. Zinc
         5. Activation energy
         6. next IE's
            1. Used to find group
            2. Topic
      5. hydrogen emission spectrum
         1. Topic
            1. Topic
         2. Topic
            1. Topic
         3. Finding the series limit graphically
            1. Topic
            2. Topic
      6. Electron Affinity
         1. Topic
         2. Group 7 elements
            1. F -328 kJ mol-1
            2. Cl -349 kJ mol-1
            3. Br -324 kJ mol-1
            4. I -295 kJ mol-1
            5. less reactive as you move down
         3. Second affinty
            1. Topic
            2. Topic
            3. Topic
      7. Atomic and Ionic radius
         1. Left: Metallic radius/covalent radius, Right:van der Waals radius Metallic because thats how it is in metals and covalent bonds. Van der Waals because of the weak attractions taking place
         2. Trends: Leaving the noble gases out, atoms get smaller as you go across a period. This happeneds because protons make it tighter whilst screening stays the same
         3. Hybridised
         4. Transition elements
         5. When filled, 4s is actually higher than 3d
         6. Ionic radius
   2. Bonding
      1. Ionic Bonding
         1. Electrons move from one atom to another
         2. Cations
            1. Anion
         3. Examples
            1. Magnesium Oxide MgO
            2. Calcium Chloride CaCl2
            3. Potassium Oxide K2O
         4. Noble gas structures
            1. Positive ions with noble gas structure
            2. Negative ions with noble gas structure
            3. Its NOT about becoming noble Instead its about the release of energy over the entire reaction
            4. It take IE for each cation and it give energy for stronger bonds So the reaction that happens will be the one where (total IE + evolved energy) is greatest
      2. Single Covalent Bonding
         1. the nucleus of both atoms attract the shared pair
            1. hydrogen
            2. hydrogen chloride
         2. complex
            1. Topic
         3. Promotion of electrons atom is in an excited state
         4. so in actual fact, the atoms will try to form as many bonded pairs as it can.
         5. Hybridisation
            1. sp3
            2. atoms merge with new molecular orbit
            3. Topic
            4. promotion happens within an energy level
            5. Topic
            6. Topic
            7. In PCl3 and 5 its more profitable for the P to promote electrons thus enabling more bonds
            8. NCl5 doesnt exist because the promotion from 2s2 would have to be to 3s which is too great
         6. sharing the GCSE "noble" view
      3. Double Covalent Bonding
         1. Topic
            1. Topic
         2. Topic
            1. Sophisticated view of the bonding in Ethane
            2. First promotion in carbon from 2s2 to 2px
            3. Next, hybridisation allows s and p to form 4 bonds
            4. Double bonds with each carbon, and 2 hydrogen single bonds
   3. Types of structure
      1. Ionic Structures
         1. Sodium Chloride lattice
            1. this is called a 6:6-co-ordinated because each atom touches 6 others
         2. 6:6 co-ordinated happens because we dont want same atoms to touch and cause repulsions
         3. Topic
            1. sometimes simplified versions are drawn
         4. The more ions surrounding each makes it more stable size ratio causes this not to happen
         5. Radius Ratio
            1. anion is >73%(cation) = 8:8
            2. 41% < anion <73% = 6:6
            3. anion < 41% = 4:4
         6. Physical Properties of NaCl
            1. High melting and boiling point
            2. More charges on ions make stronger bong
            3. smaller ions hold together better
            4. Brittle
            5. This is because they repel each other apart given some stress
            6. Soluble in water (typical of ionic solids) attractions from the loan pair in oxygen attract the positive ions which can pull them apart to form co-ordinate (dative covalent) bonds
            7. Insoluble in organic solvents
            8. Electrical behaviour
            9. When melts undergoes Electrolysis this is a Chemical change rather than physical
            10. Process:
            11. Na+ moves towards the cathode to collect an electron. returning it to Na
            12. Cl- moves towards the anode to give an electron returning it to Cl (and then 2Cl)
            13. External circuit
            14. the battery keeps filling the the cathode with electrons after they have been taken
2. Physical
3. Properties of Organic compounds
4. Organic Reaction Mechanisms
5. Inorganic
6. Instrumental analysis
7. Basic Organic