Covalent Bond & Metallic Bonding

Front 1

What is covalent bond?

Back 1

It is a bond formed by the sharing of electrons between 2 atoms

Front 2

Why do covalent substances have low melting point and boiling point?

Back 2

- within each molecule in covalent substance, the atoms are held together by strong covalent bonds

- However, molecules of covalent substances are held together by weak intermolecular forces

- A little energy is required to overcome these weak intermolecular forces. So the melting point and boiling point of covalent substances are low.

Front 3

Solubility of covalent compounds

Back 3

Soluble in organic solvent and not soluble in water

Front 4

Why do simple molecular structure does not conduct electricity?

Back 4

covalent substances do not contain mobile ion or free moving electron to carry charged particles from one end to another so covalent substances cannot conduct electricity

Front 5

Why does giant molecular structures have very high melting and boiling points?

Back 5

A giant molecular structure consists of a large number of atoms that are held together by strong covalent bonds. To melt or boil the substance, these strong bonds must be broken first.

This requires a large amount of energy.

Due to their high melting and boiling points, substances with giant molecular structures are solid at room temperature.

Front 6

Why do covalent substances in giant molecular structures are insoluble in both water and organic solvents?

Back 6

This is because in giant molecular substances, all the atoms are held together by strong covalent bonds. Any forces of attraction between the giant molecular structure and solvent molecules are not strong enough to break the strong covalent bonds.

Front 7

Why do substances with giant molecular structures do not conduct electricity (except graphite)?

Back 7

In giant molecular substances (except graphite), all the outer electrons of the atoms are used form covalent bonds. Thus, there are no free electrons that move through the structure to conduct electricity.

Front 8

Diamond is hard and has a high melting point. Explain.

Back 8

Each carbon atom is covalently bonded to 4 other carbon atoms in the shape to tetrahedral.

It is difficult to break these strong covalent bonds.

Front 9

Diamond does not conduct electricity.

Back 9

In the diamond structure, all the outer electrons of the carbon atoms are used for bonding.

Thus, there are no free electrons that move through the structure to conduct electricity.

Front 10

Graphite has a high melting point in a high boiling point. Explain.

Back 10

In the giant molecular structure of graphite, each carbon atom is covalently bonded to 3 other carbon atoms to form the shape of hexagon in layers.

A lot of energy is required to break the strong covalent bonds between the carbon atoms thus graphite has high melting and boiling points.

Front 11

Graphite conducts electricity. Explain.

Back 11

- each carbon atom has one outer electron not used to form covalent bonds

- the electron moves freely between the graphite layers

- the electron is delocalised and carries charge

Front 12

Graphite is soft and slippery. Explain.

Back 12

Each carbon atom is covalently bonded to 3 carbon atoms in the shape of hexagon in layers.

The layers of carbon atoms are held loosely by weak intermolecular forces of attraction.

These layers of carbon atoms can slide over each other easily when a force is applied

Front 13

Silicon dioxide has a high melting and boiling points. Explain.

Back 13

- each silicon atom is bonded to 4 Oxygen atoms and each oxygen atom is bonded to 2 silicon atoms. This forms a

3-dimensional structure.

- The strong Covalent bonds are difficult to break

Front 14

Silicon dioxide has a high melting and boiling points. Explain.

Back 14

- each silicon atom is bonded to 4 Oxygen atoms and each oxygen atom is bonded to 2 silicon atoms. This forms a

3-dimensional structure.

- The strong Covalent bonds are difficult to break

Front 15

Silicon dioxide does not conduct electricity. Explain.

Back 15

- in silicon dioxide structure, all the outer electrons of the oxygen and silicon atoms are used for bonding

-there are no free electrons that move through the structure to conduct electricity.

Front 16

What is a Metallic bond??

Back 16

A metallic bond is the force of attraction between positive metal ions and the 'sea of mobile electrons'.

Front 17

Why are metals very good in conducting electricity?

Back 17

While the atoms of metal are tightly packed, the outer electrons of the atoms can break away easily from the atoms. Thus, the outer electrons are able to move freely within the metal lattice.

These free-moving delocalised electrons allow the metal to conduct electricity.

Front 18

Why are metals good conductor of electricity?

Back 18

This is due to the movement of the free-moving delocalised electrons within the metal lattice.

Heat energy is transferred easily by the free-moving delocalised electrons in the structure.

Front 19

Why are metals good conductor of electricity?

Back 19

This is due to the movement of the free-moving delocalised electrons within the metal lattice.

Heat energy is transferred easily by the free-moving delocalised electrons in the structure.

Front 20

Why are metals (usually) have high density, mp and bp?

Back 20

Atoms in a metal are packed tightly in layers and are held together by strong metallic bonds.

A large amount of energy is required to break these strong metallic bonds.

This is accounts for the high densities, melting point and boiling point.

Front 21

Metals are malleable & ductile. Explain.

Back 21

This means that metals can be hammered into different shapes (malleable) and drawn into wires without breaking (ductile)

When a force is applied to a metal, the layers of metal atoms can slide over each other through the "sea of free-moving electrons"

The metallic bonding is not disrupted.