# Electrons, Bonding and structure

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• Created by: ashleigh
• Created on: 13-01-15 14:56
• Electrons, Bonding and structure
• Ionisation Energies
• NUclear charge
• Electron shielding
• Depends on;
• Example Question
• Q. State and Explain the trend in ionisation energies shown by the elements with the atomic numbers 2,10 and 18
• A. By looking at the periodic table, the elements with atomic numbers of 2,10 and 18 go down the group. as you go down atomic radius increases and there are more shells of electrons, meaning more electron shielding. This overweighs the fact there are more protons so a greater nuclear attraction, as the outer electrons experience less attractions, so ionisation energy decreases
• 4 marks
• Shells and Orbitals
• Maximum number of electrons in a shell= 2n^2
• The orbitals are S,PD and F
• Each atomic orbital can hold up to 2 electrons, but they must be opposite spins
• ^v
• Subshells
• Fill in this order; 1s 2s 2p 3s 3p 4s 3d 4p 4d 4f
• Exampls of Electron configuration; Boron has 5 electrons so would be 1S2,2S2,2P1
• Chemical Bonding
• Ionic bonding
• Between a metal and non-metal
• Transfer of electrons from metal atom to non metal atom
• postitive metal ion and neagtive non metal ion- forms electrostatic attraction
• Metallic bonding
• Electrostatic attraction of positive metal ions and negative delocalised electrons
• Covalent bonding
• Two non metals
• Sharing of electrons
• Can have multiple bonds- e.g double bonds
• Dative covalent bonding
• Where the shared pair of electrons is given from one atom only
• Lone pair- pair of electrons not involved in the bonding
• Molecular shapes
• Octahedral
• Bonding regions= 6 Lone pairs=0 bond angles= 90'
• Trigonal Planar
• Bonding regions=3 Lone pairs=0 bond angles=120'
• Tetrahedral
• Bonding regions= 4 Lone pairs= 0 bond angles= 109.5'
• Linear
• Bonding regions- 2 Lone pairs=0 Bond angles= 180'
• Pyramidal
• Bonding regions-3 Lone pairs-1 Bond angles= 107'
• Non linear
• Bonding regions= 2 Lone pairs=2 Bond angles= 104.5'
• Electronegativity and Polarity
• When a covalent bond forms bettween two atoms the same, there is equal sharing of electrons so it is said to be 100& covalent and no polar
• When a covalent bond forms between different atoms with different electronegativities the sharing is uneven and it is said to be polar, due to these differences
• Example- HCL
• Cl is more electronegative so cl has a greater attraction and so greater share of the shared electrons, meaning it becomes delta negative and H is delta positive
• This difference in charge is called a permanent dipole
• Molecules with polar bonds as non symmetrical, for molecules that are symmetrical the dipole cancels out
• Electronegativity increases up and to the right of teh period table
• Intermolecular forces and structure
• Three types of intermolecular forces
• Van der vaals forces
• Exists between all molecules. polar or non polar
• Van der vaals forces arise when there is an uneven distribution of electrons, causing a temporary dipole, this induces a dipole in a neighbouring molecule and further molecules, the induced dipoles attract each other
• Hydrogen bonds
• Permanent dipole-permanent dipole forces
• Polar molecules have permanent dipoles, which attrcat other permanent dipoles and form permant dipole- permanent dipole forces which are weak
• Water
• More dense than ice
• Ice has an open lattice with H bonds holding molecules apart, when ice melts H bonds colapse and molecules move closer together
• High melting and boiling point
• There are strong H bonds between water molecles that require extra energy to overcome