Periodicity

Question 1

State and explain how melting and boiling points change from Na and Al.

Answer 1

• ionic charge increases and number of delocalised electrons increases
• so electrostatic attraction / metallic bonding increases
• therefore there is more energy required to break the metallic bonds
• and so the melting / boiling point increases

Question 2

State and explain how melting points change across period 3, from Na to Ar.

Answer 2

• from Group 1 to Group 3, melting point increases as metallic bonding increases
• Group 4 has a high melting point as it has a giant covalent structures and there many covalent bonds to be broken
• Group 5 to Group 8 are simple covalent molecules with London forces between molecules, which are weak, and so require less energy to overcome
• so have lower melting points

Question 3

State and explain the trend in atomic radii down the group.

Answer 3

• nuclear charge increases
• shielding increases, and there are more full shells
• there is an overall decrease in nuclear attraction
• so the atomic radii increases

Question 4

State and explain the trend in atomic radii across a period.

Answer 4

• nuclear charge increases
• similar shielding
• outer electrons are more attracted to the nucleus and so there is a greater nuclear attraction
• so atomic radii decreases

Question 5

What is the first ionisation energy ?

Answer 5

the energy needed to form one mole of gaseous 1+ ions from gaseous atoms

Question 6

What is the equation for the first ionisation energy of Na ?

Answer 6

Na (g) => Na+ (g) + e-

Question 7

What are the three factors that affect the size of ionisation energy ?

Answer 7

• the atomic radius (distance from the nucleus )
• electron shielding ( number of full electron shells )
• nuclear charge ( pulling power of the nucleus )

Question 8

What is electron shielding ?

Answer 8

the number of full electron shells

Question 9

What is successive ionisation energy ?

Answer 9

a measure of the energy required to remove each electron in turn

Question 10

State and explain the trend in ionisation energy down a group.

Answer 10

• atomic radius increases and shielding increases
• so overall nuclear attraction decreases
• so 1st ionisation energy increases

Question 11

State and explain the trend in ionisation energy across a period.

Answer 11

• nuclear charge increases
• similar shielding
• atomic radius decreases
• more energy required to remove the highest energy electron, so 1st ionisation increases

Question 12

Where in the periodic table are the smallest first ionisation energies ?

Answer 12

bottom left

Question 13

Where in the periodic table are the largest first ionisation energies ?

Answer 13

top right

Question 14

Why is there a drop in ionisation energy between groups 2 and 3 ?

Answer 14

• highest energy in Gp 2 is in 2s, whereas in Gp 3 it is in 2p
• 2p is higher energy than 2s
• so less energy is required to remove an electron from 2p than 2s

Question 15

Why is there a drop in ionisation energy between Gp 15 and Gp 16 ?

Answer 15

• both Gp 15 and Gp 16 have the highest energy electron in 2p
• Gp 15 has unpaired electrons, whereas Gp 16 has paired electrons, so has more repulsion
• therefore less energy is required to remove the highest energy electron