Covalent Bonding

Question 1

Suggest why buckminsterfullerene has a much lower melting point than diamond.

Answer 1

• Weak intermolecular forces of attraction
• Not a giant structure
• No covalent bonds break

Question 2

What allows graphite to conduct electricity?

Answer 2

• Delocalized electrons that move throughout the structure and carry charge.

Question 3

Explain, in terms of its structure, why graphite can act as a lubricant.

Answer 3

• Layers slide past each other
• Weak intermolecular forces between layers

Question 4

Name the type of structure of diamond and explain, in terms of its bonding, why diamond has a high melting point

Answer 4

• Diamond has a giant covalent structure
• Covalent bonds are strong and lots of heat energy required to overcome

Question 5

Explain why diamond has a very high melting point

Answer 5

• Strong covalent bonds
• Giant covalent structure
• Require lots of heat energy to break bonds

Question 6

Fullerene has a simple molecular structure.

Explain why it has a low melting point.

Answer 6

• intermolecular forces of attraction between molecules
• Require little heat to overcome these forces

Question 7

The bonding in a hydrogen molecule is strong.

Explain why the boiling point of hydrogen is low

Answer 7

• There are weak intermolecular forces
• Requires little heat energy to overcome these forces

Question 8

Explain how the two atoms in a chlorine molecule are held together.

Answer 8

-Strong attraction between shared pair of electrons
- and nuclei of both chlorine atoms

Question 9

Hydrogen chloride gas dissolves in water to form solution A

Hydrogen chloride gas dissolves in methylbenene to form solution B

A teacher adds a piece of magnesium ribbon to each solution,

Explain why she observes effervescence with solution A but not with solution B

Answer 9

• Effervescence due to hydrogen gas
• Solution A is acidic
• Solution B is not acidic

Question 10

Suggest why the melting point of silicon dioxide is higher than the melting point of sodium chloride

Answer 10

• the (covalent) bonding in silicon dioxide is
  stronger than the (ionic) bonding in
  sodium chloride

Question 11

Explain why silicon dioxide has a high melting point

Answer 11

• There are covalent bonds that have to be broken and requires lots of energy to break these bonds

Question 12

State why carbon dioxide (CO2) is a gas at room temperature

Answer 12

• Weak intermolecular forces that require little energy to separate the molecules

Question 13

Explain why the melting point of sulfur dioxide is low.

Answer 13

• Weak intermolecular forces of attraction between molecules and require little energy to be broken

Question 14

Describe the metallic structure of molybdenum.

Answer 14

• There is a giant structure of positive ions
• Surrounded by delocalized electrons

Question 15

Explain why molybdenum is a good conductor of electricity.

Answer 15

• Delocalized electrons that flow throughout the structure

Question 16

Explain why molybdenum is malleable

Answer 16

• Layers and positive ions slide over each other

Question 17

Explain how the covalent bonds in the water molecule hold the oxygen and hydrogen atoms together.

Answer 17

• Strong electrostatic attraction between shared pair of electrons and nuclei of hydrogen and oxygen

Question 18

Why do the melting/boiling points with simple molecular structures increase, in general, with increasing relative molecular mass?

Answer 18

• The intermolecular forces increase with the size of the molecules
• So, larger molecules (i.e. molecules with greater relative molecular masses) have higher melting and boiling points.

Question 19

In diamond, each carbon is joined to ___ other carbons
covalently.

Answer 19

4

Question 20

In graphite, each carbon is covalently bonded to ___ other carbons.

Answer 20

3

Question 21

Why is graphite soft and slippery?

Answer 21

Graphite is bonded in layers and these layers can slide past each other easily because they have weak intermolecular forces in between these layers

Question 22

why is diamond hard but graphite soft?

Answer 22

• Diamond is hard because of their strong covalent bonds. Each carbon atom is covalently bonded to four other carbon atoms in a tetrahedral structure, forming a three-dimensional network.
• These strong covalent bonds require significant energy to break, making diamond extremely hard.
• Graphite is soft because of its layered structure. In graphite, each carbon atom is covalently bonded to three other carbon atoms in a flat hexagonal lattice, forming graphene layers.
• The layers are held together by weak intermolecular forces, which allow the layers to slide over each other easily making it soft.