Covalent Bonding

Question 1

What are covalent bonds?

Answer 1

When non-metal atoms share electrons with other non-metal atoms to obtain a full outer shell of electrons

Question 2

Are covalent bonds between atoms strong or weak?

Answer 2

Covalent bonds between atoms are very strong

Question 3

What is formed when two or more atoms are covalently bonded together?

Answer 3

When two or more atoms are covalently bonded together, they form ‘molecules’

Question 4

What exists between individual molecules? And are they weaker or stronger than covalent bonds?

Answer 4

Weak intermolecular forces exist between individual molecules. These forces are much weaker than covalent bonds.

Question 5

What are shared electrons called and how do they occur?

Answer 5

Shared electrons are called bonding electrons and occur in pairs

Question 6

What are electrons on the outer shell which are not involved in the covalent bond(s) called?

Answer 6

Electrons on the outer shell which are not involved in the covalent bond(s) are called non-bonding electrons

Question 7

Do simple covalent molecules electricity why or why not?

Answer 7

Simple covalent molecules do not conduct electricity as they do not contain free electrons. No free delocalised electrons

Question 8

Explain the electrostatic attraction between covalent bonds

Answer 8

There is a strong electrostatic attraction between the shared pair of electrons and the nuclei of the atoms involved, since the electrons are negatively charged and the nuclei are positively charged

Question 9

What should covalent bonds be regarded as and why?

Answer 9

Covalent bonds are best regarded as charge clouds, due to the fact that the electrons are in a state of constant motion.

Question 10

What structures do covalent bonds tend to have and give examples?

Answer 10

Covalent substances tend to have small molecular structures, such as Cl2, H2O or CO2

Question 11

Describe simple molecular structures

Answer 11

Simple molecular structures have covalent bonds joining the atoms together, but intermolecular forces that act between neighbouring molecules
They have low melting and boiling points as there are only weak intermolecular forces acting between the molecules
These forces are very weak when compared to the covalent bonds and so most small molecules are either gases or liquids at room temperature
Often the liquids are volatile

Question 12

Why do simple molecular structures have low melting and boiling points?

Answer 12

They have low melting and boiling points as there are only weak intermolecular forces acting between the molecules

Question 13

Why are most small molecules either gases or liquids at room temperature?

Answer 13

They have low melting and boiling points as there are only weak intermolecular forces acting between the molecules
These forces are very weak when compared to the covalent bonds and so most small molecules are either gases or liquids at room temperature

Question 14

What happens as the molecules increase in size?

Answer 14

As the molecules increase in size the intermolecular forces also increase as there are more electrons available
This causes the melting and boiling points to increase

Question 15

Describe what happens as the relative molecular mass of a substance increases?

Answer 15

An increase in the relative molecular mass of a substance means that there are more electrons in the structure, so there are more intermolecular forces of attraction that need to be overcome when a substance changes state
So larger amounts of heat energy are needed to overcome these forces, causing the compound to have a higher melting and boiling point

Question 16

Why are covalent compounds poor conductors of electricity?

Answer 16

They are poor conductors of electricity as there are no free ions or electrons to carry the charge
Most covalent compounds do not conduct at all in the solid state and are thus insulators

Question 17

What happens when a covalent molecule melts or boils?

Answer 17

When a covalent molecule melts or boils the covalent bonds do not break, only the intermolecular forces.

Question 18

What differs simple molecules to giant covalent structures?

Answer 18

simple molecules contain fixed numbers of atoms
Giant covalent structures on the other hand have a huge number of non-metal atoms bonded to other non-metal atoms via strong covalent bonds

Question 19

Fill in the blanks - giant lattices and have a - - of atoms in the overall structure

Answer 19

Giant lattices and have a fixed ratio of atoms in the overall structure

Question 20

Fill in the blanks - Diamond, graphite and C60 Fullerene are - of carbon

Answer 20

Diamond and graphite are allotropes of carbon

Question 21

What are allotropes?

Answer 21

Different atomic or molecular arrangements of the same element in the same physical state.

Question 22

Why aren’t diamond and graphite the same?

Answer 22

Both substances contain only carbon atoms but due to the differences in bonding arrangements they are physically completely different

Question 23

Describe the structure of a diamond

Answer 23

In diamond, each carbon atom bonds with four other carbons, forming a tetrahedron. All the covalent bonds are identical, very strong and there are no intermolecular forces

Question 24

Physics properties of diamond

Answer 24

It does not conduct electricity
It has a very high melting point
It is extremely hard and has a density of 3.51 g / cm3 – a little higher than that of aluminium

Question 25

Why is diamond very hard?

Answer 25

In diamond, each carbon atom bonds with four other carbons, forming a tetrahedron. The four covalent bonds are very strong and extend in a giant lattice, so a very large amount of heat energy is needed to break the lattice - this makes it very hard.

Question 26

What is diamond used in?

Answer 26

Diamond is used in jewellery and for coating blades in cutting tools
The cutting edges of discs used to cut bricks and concrete are tipped with diamonds
Heavy-duty drill bits and tooling equipment are also diamond tipped

Question 27

What state are giant covalent molecules in at room temperature and why?

Answer 27

Giant covalent molecules are always solids at room temperature.
•Giant covalent molecules have millions of strong covalent bonds.
They always have high melting and boiling points. This means a lot of energy is required to break them.

Question 28

Can diamonds conduct electricity why or why not?

Answer 28

•Diamond cannot conduct electricity.
• There are no free electrons to carry electrical charge.

Question 29

What is the melting point of diamond?

Answer 29

Over 3,700^oC

Question 30

Explain, in terms of electrostatic attractions, what is meant by a covalent bond.

Answer 30

The attraction between shared pairs of electrons; [1 mark]
And the nuclei of two / both atoms (in the bond)

Question 31

What is silicon (IV) oxide/silicon dioxide?

Answer 31

Silicon(IV) oxide (also known as silicon dioxide or silica), SiO2, is a giant covalent compound.

Question 32

Silicon dioxide occurs naturally as what?

Answer 32

It occurs naturally as sand and quartz

Question 33

Describe the atoms in silicon dioxide

Answer 33

Each oxygen atom forms covalent bonds with 2 silicon atoms and each silicon atom in turn forms covalent bonds with 4 oxygen atoms

Question 34

What is the form of silicon dioxide?

Answer 34

A tetrahedron is formed with one silicon atom and four oxygen atoms, similar to diamond

Question 35

Describe the forces within silicon dioxide

Answer 35

SiO2 has lots of very strong covalent bonds and no intermolecular forces so it has similar properties to diamond

Question 36

How is silicon dioxide similar to Diamond?

Answer 36

It is very hard, has a very high boiling point, is insoluble in water and does not conduct electricity

Question 37

What are some of the properties of silicon dioxide?

Answer 37

It is very hard, has a very high boiling point, is insoluble in water and does not conduct electricity

Question 38

Is silicon dioxide cheap or expensive and why?

Answer 38

SiO2 is cheap since it is available naturally and is used to make sandpaper and to line the inside of furnaces

Question 39

Describe the atoms in graphite

Answer 39

Each carbon atom in graphite is bonded to three others forming layers of hexagons, leaving one free electron per carbon atom. These free electrons migrate along the layers and are free to move and carry charge, hence graphite can conduct electricity.

Question 40

Why can graphite conduct electricity?

Answer 40

Each carbon atom in graphite is bonded to three others forming layers of hexagons, leaving one free electron per carbon atom
These free electrons migrate along the layers and are free to move and carry charge, hence graphite can conduct electricity

Question 41

Why is graphite soft and slippery?

Answer 41

The covalent bonds within the layers are very strong, but the layers are attracted to each other by weak intermolecular forces, so the layers can slide over each other making graphite soft and slippery.

Question 42

Describe the bonds in graphite

Answer 42

The covalent bonds within the layers are very strong, but the layers are attracted to each other by weak intermolecular forces, so the layers can slide over each other making graphite soft and slippery.

Question 43

What are the physical properties of graphite?

Answer 43

It conducts electricity and heat
It has a very high melting point
It is soft and slippery and less dense than diamond (2.25 g / cm3)

Question 44

What is graphite used for?

Answer 44

It is used in pencils and as an industrial lubricant, in engines and in locks
It is also used to make inert electrodes for electrolysis, which is particularly important in the extraction of metals such as aluminium

Question 45

What are fullerenes?

Answer 45

Fullerenes are a group of carbon allotropes which consist of molecules that form hollow tubes or spheres

Question 46

What are some functions of fullerenes?

Answer 46

Fullerenes can be used to trap other molecules by forming around the target molecule and capturing it, making them useful for targeted drug delivery systems
They also have a huge surface area and are useful for trapping catalyst molecules onto their surfaces making them easily accessible to reactants so catalysis can take place

Question 47

What is the structure of c60 fullerene?

Answer 47

C60 fullerene has a simple molecular structure.

Question 48

Describe the forces in c60 fullerene

Answer 48

In c60 fullerene, there are c60 molecules with weak intermolecular forces between them.

Question 49

What are the properties of c60 fullerene?

Answer 49

C60 fullerene has lower melting and boiling points than diamond and graphite.

C60 fullerene is not as hard as diamond.

C60 fullerene does not conduct electricity

Question 50

Explain why C60 Fullerene has a lower melting and boiling point than diamond and graphite.

Answer 50

When fullerene is melted, only the relatively weak intermolecular forces of attraction must be broken. This does not require as much energy as breaking all the strong covalent bonds when diamond and graphite are melted.

Question 51

Explain why C60 Fullerene is not as hard as diamond

Answer 51

It does not take as much energy to break the intermolecular forces of attraction in C60 Fullerene compared to breaking the strong covalent bonds in diamond.

Question 52

Explain why C60 Fullerene does not conduct electricity

Answer 52

Although all the carbon atoms in
C60 only form three bonds, the fourth electron on each atom can only move around within each C60 molecule; the electrons cannot jump from molecule to molecule.