Chemsitry Module 2 - Structures Of Covalent Compounds

Front 1

name the 2 types of structures that covalent compounds can have

Back 1

simple molecular lattice and giant covalent lattice

Front 2

what is a simple molecular lattice?

Back 2

3d structure of molecules bonded together by weak intermolecular fores

Front 3

give 5 examples of a simple molecular structure

Back 3

Neon, hydrogen, oxygen, nitrogen and water

Front 4

in a solid simple molecular lattice, how are the molecules and atoms held together?

Back 4

the molecules are held together by weak forces BETWEEN molecules, but the atoms within each molecule are bonded STRONGLY together by covalent bonds

Front 5

when solid iodine changes state, what happens to the Van Der Waal's forces?

Back 5

they're weak, so they break

Front 6

what are the properties of simple molecular structures?

Back 6

low melting and boiling points: intermolecular forces are weak Van Der Waal's forces, so hardly any energy at all is needed to break them

non-conductors of electricity: because there are no charged particles free to move

soluble in non-polar solvents: because van der Waal's forces form between the simple molecular structure and the non-polar solvent which in turn weakens the lattice structure

Front 7

what is a giant covalent structure?

Back 7

3d structure of atoms bonded together by strong covalent bonds

Front 8

what are 2 examples of a giant covalent structure?

Back 8

diamond and graphite

Front 9

describe the properties of giant covalent structures

Back 9

high melting and boiling points: high temps are needed to break the strong covalent bonds in the lattice

non-conductors of electricity: no charged particles are free to move (except graphite which has delocalised electrons between the layers)

insoluble in polar and non-polar solvents: because the covalent bonds are too strong to be broken by both polar and non-polar solvents.

Front 10

compare the properties of diamond and graphite

(hardness)

Back 10

hardness:
DIAMOND: = hard because the tetrahedral shape allows external forces to be spread throughout the lattice
GRAPHITE: = soft because bonding within each layer is strong and the weak forces between the layers allow the layers to slide over each other easily.

Front 11

compare the properties of diamond and graphite

(electrical conductivity)

Back 11

DIAMOND = poor conductor of electricity because there are o delocalised electrons to pick up the current of electricity as all the outer-shell electrons are used up in the covalent bonds - makes 4 bonds

GRAPHITE: = good conductor of electricity because there are delocalised electrons between layers as graphite only makes 3 bonds to carbon, so there is a free electron to move. electrons are free to move parallel to the layers when a voltage is applied

Front 12

compare the properties of diamond and graphite

(structure)

Back 12

DIAMOND: = tetrahedral structure held together by strong covalent bonds throughout the lattice

GRAPHITE: = strong hexagonal layer structure but with weak van der Waal's forces between the layers