Physical 1 Section 1 - Unit 3: Bonding

• The large difference in electronegativity between Hydrogen and the other element • Causes a dipole to be formed • So the lone pair of electrons on the other element are attracted towards the δ+H

• The power of an element to attract electrons towards itself • In a covalent bond

• CCl4 is a tetrahedral so there is rotational symmetry • This means that all the bonding pairs are facing in opposing directions • The bonds cancel out so the overall molecule is not polar

• Electron pairs repel each other and try to get to a position of maximum separation • Lone pairs repel more strongly than bonded pairs so the bond angle is higher

• Any electron pairs repel each other in molecules. • Electrons that are in a lone pair will repel other electrons more than bonding electrons. • Therefore, CCl3− will have a trigonal pyramidal shape. • And the bond angles in CCl3− will be 107°. • The shape of CCl4 will be tetrahedral. • And the bond angles in CCl4 will be 109.5°.

• There is increased nuclear charge • Electrons are added to the same shell • Therefore the outer electrons are pulled in more strongly

• Na has metallic bonding so there is attraction between the positive ions and the delocalised electrons • Na has a giant/lattice structure • There is ionic bonding in NaBr and NaI • Where there is attraction/ bonding between the oppositely charged ions in a giant lattice structure • The ionic bonds are stronger than the metallic bonds • And there is stronger attraction between the oppositely charged ions in NaBr than in NaI since the Br– ion is smaller than the I– ion

• Van der Waals • One molecule induces dipole in neighbouring molecule • Symmetrical molecule so dipoles cancel out

• Both Na+ and F have the same electron arrangement - But sodium (ion) has more protons so it attracts (outer) electrons closer

• Difference in electronegativity leads to bond polarity - And there is an attraction between ∂+ on one molecule and ∂- on another

• Bromine is (simple) molecular • Magnesium is metallic • Br2 has weak (van der Waals) forces between the molecules • So more energy is needed to overcome the stronger metallic bonds • Mg has a much greater liquid range because forces of attraction in liquid metal are strong(er)

• Electrostatic force of attraction between oppositely charged ions • In a lattice

• Shared pair of electrons

• When the shared pair of electrons in the covalent bond come from only one of the bonding atoms

• Electrostatic force of attraction between the positive metal ions and the delocalised electrons

• Number of protons (more protons, stronger bond) • Number of delocalised electrons (more electrons, stronger bond) • Size of ion (smaller ion, stronger bond)

• High melting/boiling point • Giant lattice of ions with strong electrostatic forces between oppositely charged ions • Soluble in water • Poor conductor of heat/electricity • Ions are fixed in lattice

• Low melting/boiling point • Weak intermolecular forces • Not soluble in water • Poor conductor of heat/electricity • Since there are no ions and electrons are fixed

• High melting/boiling point • Many strong covalent bonds which require a lot of energy to overcome • Insoluble in water • Diamond and sand - poor conductors since electrons can’t move • Graphite - good conductors as free delocalised electrons between layers

• High melting/boiling point • Strong electrostatic forces between positive ions and sea of delocalised electrons • Insoluble in water • Good conductors of heat/electricity • As delocalised electrons can move through structure

• 2 bp • 0 lp • 180°

• 3 bp • 0 lp • 120°

• 4 bp • 0 lp • 109.5°

• 3 bp • 1 lp • 107°