U1 Periodicity

Question 1

Order of elements in the periodic table

Answer 1

group - similar chemical properties resulting from common outer electron number

periods - increasing atomic number, demonstrated increasing number of outer electrons and a move from metallic to non-metallic characteristics

Question 2

Metallics examples

Answer 2

Li, Be, Na, Mg, Al, K, Ca

Question 3

Metallics key properties

Answer 3

1. Electrons free to move - conductive
2. Solid with high mp and bp due to closely packed lattice structure with many bonds to break.
3. More outer electrons the stronger the metallic bond.
4. Strength increases along the period because outer electrons increase
5. strength decreases down group 1 because outer elements are further from the nuclear charge

Question 4

Covalent molecular examples

Answer 4

H2, N2, O2, F2, Cl2, Br2, I2, P4, S8 C60

Question 5

Covalent molecular key properties

Answer 5

1. In all groups 5,6 and 7, molecular elements the intramolecular forces are all strong covalent bonds. 2. The intermolecular forces are weak London Dispersion Forces.
2. Moving down group bp increases because the strength of London dispersion forces increases.

Question 6

Covalent Network examples

Answer 6

B, C (diamond/graphite), Si

Question 7

Covalent network properties

Answer 7

1. High melting and boiling points because a lot of energy is needed to break strong covalent bonds.

Question 8

Monatomic elements examples

Answer 8

He, Ne, Ar (noble gases)

Question 9

Monatomic elements

Answer 9

Full outer electron shell, therefore don’t bond with other atoms, and no free electrons so no conductivity.

London dispersion forces between monatomic elements

Question 10

Covalent radius

Answer 10

Measure of size of atoms.

Defined as half the distance between centres (nuclei) of 2 bonded atoms. To find bond length add 2 covalent radii together.

Trends in covalent radius across periods and down groups can be explained in terms of occupied shells and nuclear charge.

Question 11

Trends in covalent radius

Answer 11

Across the period, nuclear charge and the number of outer elctrons increase and attract the outer electrons closer to the nucleus - atomic size decreases

Down a group, electron shells increase but the number of outer electrons stays the same, so atomic size increases.

Question 12

First ionisation energy

Answer 12

The amount of energy required to remove one mole of electrons from one mole of atoms in gaseous state

E(g) arrow > E+(g) + e-

The outer most electron will be the most weakly held and is removed first. The ionisation energy is an enthalpy change and therefore is measured per mole - KJmol-1

Question 13

Second Ionisation Energy

Answer 13

Defined as “the amount of energy required to remove one mole of electrons from one mole of gaseous 1+ ions”

E+(g) arrow> E2+(g) + e-

Question 14

3 key parts to ionisation energy definitions

Answer 14

1. atoms must be in gaseous state
2. one mole of atoms is compared
3. one mole of electrons is removed

Question 15

Trends in ionisation energy across period

Answer 15

Across the period, nuclear charge increases (greater positive charge on the nucleus) and holds the outer electrons more strongly. More energy needs to be supplied to remove the electron. Therefore ionisation energy energy increases.

Question 16

Trends in ionisation energy down a group

Answer 16

Electrons further away from the nucleus

The screening effect of inner electron shells reduces the nuclear attraction for the outer electrons, despite the increased (pos) nuclear charge. Therefore, ionisation energy decreases.

Question 17

Electronegativity

Answer 17

Atoms of different elements have different attractions for bonding electrons.

Electronegativity is a measure of the attraction of an atom involved in a bond has for the electrons of the bond. Values in data booklet

Question 18

Trends in electronegativity across period

Answer 18

Nuclear charge increases, attracting the electrons strongly to the nucleus. As a result, electronegativity increases.

Question 19

Trends in electronegativity down a group

Answer 19

The nuclear charge increases but the number of electron shells also increases. As a result of ‘shielding’ and the inner distance the outer shell is from the nucleus, the electronegativity decreases.