CHAPTER 7: PERIODICITY

Question 1

What is the structure of diamond?

Answer 1

In diamond, each carbon atom is covalently bonded to four other carbon atoms.

Question 2

What are the properties of diamond?

Answer 2

• High melting point (>3800 K) — diamond has incredibly strong covalent carbon-carbon bonds that require lots of energy to break
• Hard — diamond is one of the hardest known materials due to its incredibly strong covalent bonds
• Electrical insulator — all of carbon’s electrons are involved in covalent bonding so no electrons are free to move, meaning there are no mobile charge carriers and hence that it cannot conduct electricity
• Insoluble — diamond is insoluble in all solvents because the covalent bonds in diamond are stronger than any possible forces of attraction with any solvent
• Thermal conductor — vibrations can easily travel through it, making it good at conducting heat

Question 3

What is the structure of graphite?

Answer 3

Graphite has a layered structure. In each layer, every carbon atom is covalently bonded to three other carbon atoms.

Each carbon has a free electron held in a p-orbital. These p-orbitals can overlap to
produce a ‘cloud’ of electrons above and below each layer; this is delocalised and can carry current.

Question 4

What are the properies of graphite?

Answer 4

• High melting point — graphite has incredibly strong covalent carbon-carbon
bonds that require lots of energy to break

• Soft and slippery — the layers can easily slide over each other due to the weak London forces of attraction between them – there is no covalent bonding between the layers
• Low density — less dense than diamond because of the distance between layers
• Electrical conductor — delocalised electrons that carry current are free to move between the layers, acting as mobile charge carriers. Graphite only conducts electricity in hexagonal planes, and not at right angles to these plane
• Insoluble — graphite is insoluble in all solvents because the covalent bonds in graphite are stronger than any possible forces of attraction with any solvent

Question 5

What is the structure of graphene?

Answer 5

Graphene is a sheet of carbon atoms that is only one atom thick. The carbon atoms are joined together to form hexagons, which are held in strong covalent bonds.

Question 6

What are the properties of graphene?

Answer 6

• The delocalised electrons are free to move, so it is a great electrical conductor
• A layer of graphite is very light — it is the thinnest material ever made
• The covalent bonds are strengthened by the delocalised electrons and so a single layer of graphene is very strong
• High melting point and insoluble in all solvents for the same reasons as diamond and graphite

Question 7

What is the definition of a metallic bond?

Answer 7

The strong electrostatic attraction between positive metal ions (cations) and negative delocalised electrons in a metal lattice.

Question 8

Explain electrical conductivity in metals.

Answer 8

• Metals are incredibly good at conducting electricity because of their delocalised electrons
• They conduct electricity in both the solid and liquid states

• The delocalised electrons are mobile charge carriers which are free to move
throughout the lattice and carry current, allowing metals to conduct electricity

Question 9

Explain solubility in metals.

Answer 9

• Metals do not dissolve like ionic compounds, because any interactions between a polar solvent and the charges in the lattice would lead to a reaction rather than dissolving

Question 10

Explain melting and boiling points in metals.

Answer 10

• A lot of energy is required to overcome the strong metallic bonds. Melting and boiling points are therefore very high
• The melting point is influenced by several factors, including the charge and size of the metal ion.
• Metal ions of greater charge contribute more electrons to the delocalised sea, therefore there are more electrostatic forces of attraction between the ions and electrons, resulting in a higher melting point
• The positive nuclei of smaller cations are closer to the delocalised electrons, resulting in a stronger metallic bond, and therefore a higher melting point

Question 11

Explain malleability and ductility in metals.

Answer 11

• There are no bonds holding specific ions together — this allows ions to slide and move past each other
• Metals are malleable and ductile — they can be hammered into different shapes (malleable) or drawn out and stretched (ductile). This is because the delocalised electrons can move. This means the structure can ‘give’ slightly, allowing atoms to slide past each other

Question 12

What is first ionisation energy?

Answer 12

The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.
X (g) - X+ (g) + e-

Question 13

How does atomic radius affect ionisation energy?

Answer 13

The larger the atomic radius, the further away the outer electrons are from the nucleus, and the less nuclear attraction experienced by the outer electrons. Therefore, it will be easier to remove an electron and the ionisation energy will be lower.

Question 14

How does nucleur charge after ionisation energy?

Answer 14

The greater the nuclear charge, the greater the attractive force experienced by the outer electrons, and the harder it will be to remove an electron. The ionisation energy will be greater.

Question 15

How does shielding affect ionisation energy?

Answer 15

Electrons are negatively charged and will tend to repel each other. The greater the number of inner shells of electrons, the greater the repulsion of the outer shell of electrons and the easier it will be to remove an electron. The ionisation energy will be lower.

Question 16

What is the trend in ionisation energy across a period?

Answer 16

• Ionisation energy increases across a period
• The number of protons in the nucleus and the number of electrons in the outer shell increases, but the electrons are all added to the same shell
• There is a greater attraction between the nucleus and the outer electrons, so it becomes harder to remove an electron when moving across a period

Question 17

What is the trend in ionisation energy down a group?

Answer 17

• Ionisation energy decreases down
  a group
• The number of shells increases so the atomic radius and shielding increase
• There is less attraction between the nucleus and the outer electrons, so it becomes easier to remove an electron down when moving down a group

Question 18

What happens from Group 1 to Group 4?

Answer 18

From Group 1 to Group 14 (4), the melting point increases steadily — these elements all have giant structures:
• Giant Metallic Structures — as the nuclear charge increases, the number of delocalised electrons also increases. This allows for stronger attractions between cations and electrons, so more energy is needed to overcome the stronger bonds across the period
• Giant Covalent Structures — each successive element has an additional electron, allowing more covalent bonds to be formed. This means that more energy is needed to overcome strong covalent bonds

Question 19

What happens from Group 4 to Group 5?

Answer 19

From Group 14 (4) to Group 15 (5), the melting point sharply decreases — there is a change in structure:
• The structure changes from giant to simple molecular, so weak intermolecular forces exist between atoms. These are London forces of attraction. These forces need little energy to overcome which means the melting point decreases significantly from Group 14 (4) to Group 15 (5)

Question 20

What happens from Group 5 to Group 8?

Answer 20

From Group 15 (5) to Group 18 (8), the melting point remains relatively low — these elements are all simple molecular:
• Simple molecular structures exist with only weak intermolecular forces (London forces) of attraction between molecules, so melting point stays low
• Noble gases are monoatomic, so their molecules have the fewest electrons and therefore the weakest London forces of attraction of the simple molecular substances

Question 21

What are the properties of Group 2?

Answer 21

Group 2 elements have similar physical properties:
• Relatively high melting and boiling points — these decrease down the group due to the increased atomic radii and shielding as the atom gets larger — the electrostatic attraction between the delocalised electrons and the nucleus of the ions decreases and so the metallic bonding is weaker
• Low density, light metals
• Form white compounds

Question 22

What happens when Group 2 elements react with oxygen?

Answer 22

Group 2 elements react vigorously with oxygen to form an ionic oxide with the general formula MO.

Question 23

What happens when Group 2 elements react with water?

Answer 23

Group 2 elements, excluding beryllium, react with water to form metal hydroxides with the general formula M(OH)2 and hydrogen gas. The reactions with cold water become increasingly vigorous moving down the group, with barium reacting most vigorously.

Question 24

What happens when Group 2 elements react with dilute acids?

Answer 24

Group 2 elements, excluding beryllium, react with dilute acids to form a salt and
hydrogen gas. The reaction becomes more vigorous moving down the group. Different salts are produced depending on the acid used.

Question 25

How can magnesium compounds be used?

Answer 25

• Indigestion remedies — Mg(OH)2 (milk of magnesia) is used to neutralise excess
hydrochloric acid in the stomach, acting as an antacid, relieving indigestion

Question 26

How can calcium compounds be used?

Answer 26

• Agriculture — Ca(OH)2 (hydrated lime) is used to neutralise acidic soil
• Building materials — limestone and marble are used in construction and contain CaCO3. However, CaCO3 is basic. This means that buildings containing CaCO3 are susceptible to erosion by acid rain
• CaCO3 is also used as an indigestion remedy (an antacid)