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Electrons, bonding and structure

Electron configuration
1. Orbitals; contain two electrons in orbit around the nucleus
  1. 1st energy level; 1s
  2. 2nd energy level; 2s and 3 lots of 2p
  3. 3rd energy level; 3s, 3 lots of 3p and 5 lots of 3d
  4. 4th energy level; 4s, 3 lots of 4p, 5 lots of 4d and 7 lots of 4f
  5. 4s will fill up before 3d
  6. orbitals will fill up s,p,d,f
2. Rules of electron configuration and orbitals
  1. Aufbau's principle; electrons fill up the lowest possible energy levels first (they are attracted to the nucleus
  2. Hund's rule of multiplicity; orbitls are occupied singularly before they begin to pair up (they repel each other due to negative charges)
  3. Spin pairing; a pair of electrons will spin in opposite directions when in an orbital
3. orbital shapes
  1. S orbital
    1. a circle
  2. P orbital
    1. a figure of 8
Bonding
1. Metallic bonding; the electrostatic attraction between positive metal ions and delocalised electrons
  1. positive ions lined up next to each other with delocalised electrons
  2. metal and a metal
2. Covalent bonding; the electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms
  1. non-metal and a non-metal ion
  2. dot and cross diagram
  3. Covalent bonding exceptions; BF3 and SF6
3. Ionic bonding; the electrostatic attraction between oppositely charged ions
  1. non-metal and a metal
  2. each atom in a box, charges indicated, dot and cross diagram in boxes
4. Dative covalent bonding; a covalent bond where both electrons come from the same atom
  1. dot and cross in a box with charge
Molecule shape
1. Two bonding pairs
  1. separate by 180`
  2. linear
  3. BeCl2
  4. 02
2. Three bonding pairs
  1. separate by 120`
  2. BF3
  3. trigonal planar
3. Four bonding pairs
  1. separate by 109.5`
  2. CH4
  3. tetrahedral
4. Five bonding pairs
  1. separate by 90` at the poles and 120` at the equator
  2. PCl5
  3. trigonal bipyramidal
5. Six bonding pairs
  1. separate by 90`
  2. octahedral
  3. SF6
6. Lone pairs
  1. reduce the bond angle by 2.5` (per lone pair)
  2. NH3
  3. trigonal pyramid
7. Water
  1. non-linear
  2. has two lone pair of electrons
Intermolecular forces
1. London forces (i.d - i.d)
  1. occur in non-polar molecules and hydrocarbons
  2. instantaneous dipole; i.d molecule will force a neighbouring molecule to form an induced dipole
  3. more electrons means a greater likelihood that the molecules will become instantaneously dipole (F2, Cl2, Br2, I2)
2. Permanent dipole - induced dipole (p.d - p.d)
  1. molecules with a permanent dipole will be attracted to each other
  2. occurs due to polar bonding
3. Permanent dipole - permanent dipole (p.d - p.d)
  1. occurs due to polar bonding
  2. if a polar molecule has a permanent dipole it can induce a dipole in a neighbouring molecule
    1. nitrogen, oxygen, fluorine, bromine, chlorine, iodine and sulphur cause a permanent dipole
4. Hydrogen Bonds
  1. the attraction between an electron deficient hydrogen and the lone pair of and electronegative atom in another molecule
  2. Anomalous properties of water
    1. Relatively high melting and boiling point; lots of energy is required to break the relatively strong hydrogen bonds.
    2. Ice floats; the molecules of ice are held apart in an open lattice by hydrogen bonds
5. Dipole; the uneven distribution of electrons in a covalent bond
6. Electronegativity; the ability of an atom to attract the electron pair in a covalent bond to itself
  1. nitrogen, oxygen, fluorine, bromine, chlorine, iodine and sulphur

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Electrons, bonding and structure

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	Created by Raisha Gibbs almost 8 years ago