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• Standard Enthalpy Change of Reaction

  Enthalpy change that accompanies a reaction in the molar quantities expressed in the chemical equation under standard condition.

  
  

  (Total BE reactants) - (Total BE products)
• Standard Enthalpy change of Combustion

  When 1 mole of a substance reacts completely with oxygen under standard conditions.

  
  

  1. q = mc△t/1000 -> m = Mass of H2O
  2. Moles of fuel
  3. kJ/Moles = kJ mol^-1
• Standard Enthalpy change of Neutralisation

  When 1 mole of water is formed in an acid-base reaction under standard condition.

  
  

  1. q = mc△t/1000 -> m = Vol. of acid + base
  2. Moles
  3. kJ/Moles = kJ mol^-1
• Standard Enthalpy change of Formation

  When 1 mole of a compound is formed from its elements under standard conditions.

  
  

  1. q = mc△t/1000
  2. Moles of limiting reactant
  3. kJ/Moles = kJ mol^-1
• Relative Atomic Mass

  Weighted mean mass of of one atom compared to one 12th of the mass of 1 atom of Carbon-12
• Empirical Formula

  Simplest ratio of atoms of each element in the compound
• Oxidation

  Losing electrons, losing hydrogen -> Positive -> Increases oxidation number
• Reduction

  Gaining electrons, gaining hydrogen -> negative -> decreases oxidation number
• Rules for Oxidation numbers

  1. Uncombined elements = 0
  2. Sum of numbers in a compound = 0
  3. Monoatomic ions = charge
  4. Polyatomic ions = sum of OZ
• Invariable ON

  Group 1 = +1

  Group 2 = +2

  Al = +3

  H = +1 -> non metal and -1 -> metals

  F = -1

  Cl, Br, I = -1

  O = -2, -1 w/ F
• Disproportionation

  When the same species gets reduced and oxidised.
• Electron Structure

  Principal Level (1,2,3,4...) -> Sub levels (spdf) -> Orbitals (Holds up to 2 e- in opposite spins)

  
• Ionic bonding

  Electrostatic force of attraction between 2 oppositely charged ions.
• Ionic bonding

  Ionic bonding is stronger and the melting points higher when the ions are smaller and/ or have higher charges.

  Giant ionic lattice

  High MP

  Non conductors when solid = Ions aren't free to move

  Usually soluble in aqueous solutions
• Covalent Bonding

  Strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms

  Simple Molecular
• Dative Covalent Bonding

  When the shared pair of e- in the covalent bond comes from only one of the bonding atoms.
• Giant Ionic

  Mp + Bp = High -> Giant lattice of ions with strong electrostatic forces

  Good solubility in water

  Poor conductors when solid -> Ions can't move

  Good conductors when liquid -> ions are free to move
• Simple Molecular

  Mp + Bp = Low -> Weak intermolecular forces b/t molecule -> induced dipole interaction + hydrogen bonding

  Poor solubility in water

  Poor solid conductor -> e- are fixed in place

  Poor liquid conductors -> No ions to carry charge
• Linear

  Bonding = 2

  Lone = 0

  Bond Angle = 180

  CO2, CS2, HCN, BeF2

  
• Trigonal Planar

  Bonding = 3

  Lone = 0

  Bond Angle = 120

  BF3, AlCl3, SO3, NO3-, CO3 2-
• Tetrahedral

  Bonding = 4

  Lone = 0

  Bond Angle = 109.5

  SiCl4, SO4 2-, ClO4-, H3O+

  
• Trigonal Pyramidal

  Bonding = 3

  Lone = 1

  Bond Angle = 107

  NCl3, PF3, ClO3, H3O+

  
• Bent/Non-linear

  Bonding = 2

  Lone = 2

  Bond Angle = 104.5

  OCL2, H2S, OF2, SCl2

  
• Octahedral

  Bonding = 6

  Lone = 0

  Bond Angle = 90

  SF6

  
• Explaining bond angles

  1. How many bond pairs
  2. How many lone pairs
  3. lp = 0, bond pairs repel equally and move as far apart as possible
  4. lp > 0, lp repel more than bp
• Electronegativity

  The ability of an atom to attract the bonding electrons in a covalent bond towards itself.
• Factors effecting electronegativity

  increases across a period -> no. of p+ increases = AR decreases b/c e- in the same shell are pulled in more (greater nuclear attraction)

  Decreases down a group -> distance b/t nucleus + outer e- increases + shielding
• Permanent Dipole

  Forms when elements have different electronegativities, unequal distribution of electrons in the bond creates a dipole. Molecule isn't symmetrical (all bonds identical + no lone pair)-> dipoles don't cancel.

  
  

  Long alkanes = large SA of contact b/t molecules to induce dipole-dipole interactions
• Induced dipole - dipole = London Forces


  Occur b/t all molecular substances and noble gases, not in ionic substances. Uneven distribution of e- creates a slightly -ve and slightly +ve charges on either side of the atom. When a temporary dipole is establishes, it induces a temporary dipole on its neighbouring molecules
• Hydrogen Bonding

  Occurs in compounds when Hydrogen is attached to N/O/F (available lone pair)

  Large electronegative difference between H and N/O/F.

  Strongest bond

  Type of permanent dipole dipole
• Periodicity


  A repeating pattern across different periods
• Atomic Radius

  Decreases = Left to Right -> increase no of p+ = more +ve charge attraction for e- in the same shell
• First ionisation energy

  Energy needed to remove one mole of e- from one mole of gaseous atoms to create one mole of gaseous 1+ ions.

  H(g) -.> H+(g) + e-
• Factors affecting Ionisation energy

  Nuclear Attraction -> More p+ in nucleus = greater attraction

  Atomic Radius -> Bigger atom = e- on outer shell are further away from nucleus = weaker attraction to nucleus

  Shielding -> e- in an outer shell is repelled by e- in complete inner shells (weakens nuclear attractions)
• Metallic Bonding

  Electrostatic forces of attraction between positive metal ions and delocalised electrons
• Longer straight chain = induced dipole dipole interaction

  More IDDI Forces = more energy is needed to overcome these forces (cuz more e-)

  Less IDDI Forces = Less energy is needed to overcome these forces -> since branched, they can't pack closely together which decreases surface contact b/t molecules
• Catalyst

  A substance that increases the rate of reaction by providing the reaction with an alternative pathway that has a lower activation energy. Catalyst is chemically unchanged at the end of the reaction.
• Free Radical Substitution

  Alkane -> Halagenoalkane

  
• Sigma Bond

  When 2 s orbitals over lap
• Pi bond

  Sideways overlap of a 2 p orbitals