Covalent Bonding GCSE flashcards

define what a covalent bond is

a shared pair of electrons between 2 atoms- each negatively charged electron is attracted by an electrostatic force of attraction to the adjacent atoms positively charges nucleus

example of simple molecular structures

• H₂O

• I₂

• C₆₀

simple molecular structures have ……………. between the …….. but also have ………….. that act between neighbouring ……….

• strong covalent bonds

• atoms

• weak intermolecular forces

• molecules

why do simple molecular structures have low melting and boiling points

they have weak intermolecular forces between the molecules, so little energy is needed to overcome them

what happens to the melting and boilings points of a substance as the molecule increases in size

intermolecular forces get stronger, as there are more electrons available, so melting and boiling points increase as more energy is needed to overcome the forces

draw a dot cross diagram for the simple molecules:

1. H₂O

2. NH₃ ammonia

3. CH₄ methane

4. Chloromethane CH₃Cl

5. Dichloroethane C₂H₄Cl₂

okay

how can you tell whether a compound is covalent (simple molecular) or ionic

• covalent ( simple molecular )- non metal + non metal

• ionic-e metal + non metal

explain why sodium bromide NaBr has a higher melting point than hydrogen bromide HBr

• sodium bromide is a giant ionic lattice structure

• ions are held together with strong electrostatic forces of attraction between oppositely charged ions

• means lots of energy is needed to overcome these bonds

• hydrogen bromide is a simple molecular structure

• has weak intermolecular forces between the molecules

• means little energy is needed to overcome these forces

explain how the covalent bonds in a water molecule hold the hydrogen and oxygen atoms together

• hydrogen and oxygen atoms each share a pair of electrons which are negatively charged

• the shared pair of electrons attract to the positively charges nuclei of the adjacent atom

• with strong electrostatic force of attraction

• that hold the atoms together

why does fullerene have a much lower melting point than diamond even though they are both made of carbon atoms

• C₆₀ is a simple molecular structure, where as diamonds is a giant covalent

• when melting C₆₀, weak intermolecular forces are broken, where as in diamond, strong covalent bonds are broken

describe the forces of attraction in a covalent bond

• very strong electrostatic forces of attraction between a shared pair of negatively charges electrons and positively charges nuclei of the adjacent atom

explain why C₄H₁₀ has a higher boiling point than C₂H₆

• C₄H₁₀ and C₂H₆ are both simple molecular structures with weak intermolecular forces acting between the molecules

• however the IMF’s are stronger between C₄H₁₀ as it is bigger than the C₂H₆ molecule and more electrons are available

• this means that more energy is needed to overcome the stronger forces in C₄H₁₀ and less energy is needed to over the weaker IMF’s in C₂H₆

give 3 examples of giant covalent lattice structures

• diamond

• graphite

• silicon dioxide

giant covalent lattice structures have … numbers of non metal …… bonded to other …… atoms via ………. ………. bonds

1. large

2. atoms

3. non metal

4. strong

5. covalent

why do giant covalent lattices have high melting and boiling points

• they have many strong covalent bonds between their atoms

• which need large amounts of energy in order to overcome

• hence why they are solids at room temp

what type of structure are diamond, graphite and silicon dioxide

what type of atom is diamond made from

carbon

what type of atom is graphite made from

carbon

what types of atoms is silicon dioxide made from

silicon and oxygen

what number of bonds does each carbon atom form to made diamond

each carbon atoms forms 4 bonds to neighbouring carbon atoms

what number of bonds does each carbon atom form to make graphite

each carbon atom from 3 bonds to neighbouring carbon atoms

in silicone dioxide, each silicon atom is bonded to … oxygen atoms

each oxygen atom is bonded to … silicon atoms

4

2

does diamond contain delocalised electrons

no- so it doesn’t conduct electricity

does graphite contain delocalised electrons

yes- contains 4 electrons in outer shell so can share 4, but only shares 3; so has 1 delocalised electron per carbon atom

so it can conduct electricity

does silicon dioxide contain delocalised electrons

no- so it doesn’t conduct electricity

learn how to label a diagram of silicon dioxide

okay

why does diamond have high melting and boiling points

• many strong covalent bonds between atoms

• which need large amounts of energy in order to overcome

why does graphite have a high melting and boiling point

• many strong covalent bonds between atoms

• which need large amounts of energy in order to overcome

why does silicon dioxide have high melting and boiling points

• many strong covalent bonds between atoms

• which need large amounts of energy to overcome

does diamond conduct electricity

no- does not contain any charges particles ( no ions or delocalised electrons )

so no flow of charge throughout whole structure

does graphite conduct electricity

yes- contains delocalised electrons which are free to move

so there is a flow of charge throughout the whole structure

does silicon dioxide conduct electricity

no- does not contain any charges particles ( no ions or delocalised electrons )

so no flow of charge throughout whole structure

is diamond hard or soft, and why

hard- strong covalent bonds hold atoms in fixed positions

so there is no movement within giant lattice structure

is graphite hard or soft, and why

soft- atoms form layers

layers can slide over each other

as they only have weak intermolecular force between them

is silicon dioxide hard or soft, and why

hard ish- strong covalent bonds hold atoms in fixed positions

so there is no movement within giant lattice structure

explain why fullerene has a much lower melting point than diamond and graphite

• fullerene is a simple molecular structure

• when melting, it is the weak intermolecular forces that are overcome

• diamond and graphite are giant covalent lattices

• when melting them, it is the strong covalent bonds that are overcome

• more energy is needed to overcome the strong covalent bonds than the weak intermolecular forces, hence why diamond and graphite have a higher melting point than fullerene

discuss the differences between diamond and graphite

refer to structure and bonding, electrical conductivity and hardness in your answer

• diamond and graphite are both giant covalent lattice structures, with strong covalent bonds between their atoms

• however the carbon atoms in graphite form layer that can slide over each as the layers only have weak intermolecular forces between them

• this means that graphite is soft compared to diamond, which is very hard due to its strong covalent bonds that hold the atoms in fixed positions, so there is no movement within diamonds giant lattice

• graphite can conduct electricity as it contains delocalised electrons which are free to move, so their is a flow of charge throughout the whole structure

• however diamond does not conduct electricity as it does not contain any charges particles (delocalised electrons) so there is no flow of charge throughout the whole structure

nitrogen dioxide and silicon dioxide both contain covalent bonds, explain why nitrogen dioxide has a much lower melting point than silicon dioxide

• Nitrogen dioxide is a simple molecular structure

• it has strong covalent bonds between its atoms however it also has weak intermolecular forces between its neighbouring molecules

• these intermolecular forces are overcome when nitrogen dioxide melts (not the covalent bonds) and as the intermolecular forces are weak little energy is needed to overcome them

• Silicon dioxide is a giant covalent lattice structure

• it also has a very strong covalent bonds between its atoms and (has no intermolecular forces) so these covalent bonds need to be overcome in order for silicon dioxide to melt

• this takes a lot of energy to overcome the covalent bonds hence why it has a very high melting point

Explain why magnesium sulphate conducts electricity as a liquid but not as solid

When molten the ions are free to move so there is a flow of charged particles throughout the structure, however when solid the ions are not free to move and are in fixed positions in the lattice so there is no flow charge throughout the structure

Explain why carbon dioxide has a low boiling point

It is a simple molecular structure – weak intermolecular forces between its molecules that are easily overcome by little energy