Electrons, bonding and structure

Electrons, bonding and structure

1. Electron configuration

   • Orbitals; contain two electrons in orbit around the nucleus
     • 1st energy level; 1s
     • 2nd energy level; 2s and 3 lots of 2p
     • 3rd energy level; 3s, 3 lots of 3p and 5 lots of 3d
     • 4th energy level; 4s, 3 lots of 4p, 5 lots of 4d and 7 lots of 4f
     • 4s will fill up before 3d
     • orbitals will fill up s,p,d,f
   • Rules of electron configuration and orbitals
     • Aufbau's principle; electrons fill up the lowest possible energy levels first (they are attracted to the nucleus
     • Hund's rule of multiplicity; orbitls are occupied singularly before they begin to pair up (they repel each other due to negative charges)
     • Spin pairing; a pair of electrons will spin in opposite directions when in an orbital
   • orbital shapes
     • S orbital
       • a circle
     • P orbital
       • a figure of 8

2. Bonding

   • Metallic bonding; the electrostatic attraction between positive metal ions and delocalised electrons
     • positive ions lined up next to each other with delocalised electrons
     • metal and a metal
   • Covalent bonding; the electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms
     • non-metal and a non-metal ion
     • dot and cross diagram
     • Covalent bonding exceptions; BF3 and SF6
   • Ionic bonding; the electrostatic attraction between oppositely charged ions
     • non-metal and a metal
     • each atom in a box, charges indicated, dot and cross diagram in boxes
   • Dative covalent bonding; a covalent bond where both electrons come from the same atom
     • dot and cross in a box with charge

3. Molecule shape

   • Two bonding pairs
     • separate by 180`
     • linear
     • BeCl2
     • 02
   • Three bonding pairs
     • separate by 120`
     • BF3
     • trigonal planar
   • Four bonding pairs
     • separate by 109.5`
     • CH4
     • tetrahedral
   • Five bonding pairs
     • separate by 90` at the poles and 120` at the equator
     • PCl5
     • trigonal bipyramidal
   • Six bonding pairs
     • separate by 90`
     • octahedral
     • SF6
   • Lone pairs
     • reduce the bond angle by 2.5` (per lone pair)
     • NH3
     • trigonal pyramid
   • Water
     • non-linear
     • has two lone pair of electrons

4. Intermolecular forces

   • London forces (i.d - i.d)
     • occur in non-polar molecules and hydrocarbons
     • instantaneous dipole; i.d molecule will force a neighbouring molecule to form an induced dipole
     • more electrons means a greater likelihood that the molecules will become instantaneously dipole (F2, Cl2, Br2, I2)
   • Permanent dipole - induced dipole (p.d - p.d)
     • molecules with a permanent dipole will be attracted to each other
     • occurs due to polar bonding
   • Permanent dipole - permanent dipole (p.d - p.d)
     • occurs due to polar bonding
     • if a polar molecule has a permanent dipole it can induce a dipole in a neighbouring molecule
       • nitrogen, oxygen, fluorine, bromine, chlorine, iodine and sulphur cause a permanent dipole
   • Hydrogen Bonds
     • the attraction between an electron deficient hydrogen and the lone pair of and electronegative atom in another molecule
     • Anomalous properties of water
       • Relatively high melting and boiling point; lots of energy is required to break the relatively strong hydrogen bonds.
       • Ice floats; the molecules of ice are held apart in an open lattice by hydrogen bonds
   • Dipole; the uneven distribution of electrons in a covalent bond
   • Electronegativity; the ability of an atom to attract the electron pair in a covalent bond to itself
     • nitrogen, oxygen, fluorine, bromine, chlorine, iodine and sulphur