3.1 - Periodic table

Question 1

How are elements arranged in the periodic table?

Answer 1

• increasing atomic number
• periods = repeating trends in chemical and physical properties
• groups = similar chemical properties

Question 2

Give the groups/elements which fall into s, p and d blocks and why are they classified as these?

Answer 2

s: group 1 and 2
p: right side (metals and non-metals)
d: transition metals

highest energy electron is in a s/p/d orbital

Question 3

Define first ionisation energy

Answer 3

removal of 1 mol of electrons from 1 mol of gaseous atoms, forming 1 mol gaseous 1+ ions

Question 4

Ionisation energy across a period

Answer 4

• nuclear charge increases
• same shielding
• atomic radius decreases
• overall nuclear attraction increases
• so IE increases

Question 5

Ionisation energy down a group

Answer 5

• nuclear charge increases
• shielding increases
• atomic radius increases
• overall nuclear attraction decreases
• so IE decreases

Question 6

Define metallic bonding

Answer 6

strong electrostatic attraction between fixed + cations and mobile delocalised electrons

Question 7

Properties of metals

Answer 7

• high MP: electrostatic attraction requires lot of energy to overcome
• can conduct electricity: electrons can move when voltage applied
• insoluble: no H bonds

Question 8

Bond angles and shape of diamond and silicon giant covalent lattice structures

Answer 8

109.5°
tetrahedral

Question 9

Properties of giant covalent lattices

Answer 9

• high MP/BP: strong covalent bonds
• mostly insoluble: bond strength
• can’t conduct electricity: no delocalised electrons (but graphene/graphite CAN)

Question 10

Bond angles in graphene and graphite

Answer 10

120°

Question 11

What is a solid giant covalent lattice?

Answer 11

networks of atoms bonded by strong covalent bonds

Question 12

Trend across period 3

Answer 12

• atomic radius decreases
• MP increases

Question 13

Do giant covalent lattices have London forces?

Answer 13

no, only covalent bonds

Question 14

Why do successive IEs increase with ionisation number?

Answer 14

• atomic radius decreases
• nuclear attraction increases

Question 15

Why are there some small decreases in 1st IE in graph against atomic number?

Answer 15

• outermost electron in higher energy sub-shell (further from nucleus)
• removing electron from orbital full with 2 electrons: due to repulsion, less energy required to remove one than from an orbital with only 1 electron