Periodicity

Question 1

define periodicity

Answer 1

repeating trend in physical and chemical properties across the periods of the periodic table

Question 2

the periodic table is arranged in order of increasing…

Answer 2

atomic (proton) number

Question 3

why do elements in the same group have similar physical and chemical properties

Answer 3

similar outer shell configuration

Question 4

across a period what happens to the atomic number and what does it cause

Answer 4

• atomic number increases (more protons) but same no of shells
• radius gets smaller, increasing electrostatic attraction and electronegativity

Question 5

why does IE1 decrease down a group

Answer 5

• increased SHIELDING from the nucleus so decreased attraction between the nucleus and the electron shells

Question 6

why does IE1 generally increase across a period

Answer 6

• increased nuclear charge without changing the no of shells so increased attraction between the nucleus and outer e- shells

Question 7

Exceptions to increasing IE1 across period and why (Beryllium > Boron) (Nitrogen > Oxygen )

Answer 7

• IE1 of Boron is lower than that of Beryllium because e- in boron is removed from 2p subshell which is higher in energy so less energy is required to remove it
• IE1 of Oxygen is lower than that of Nitrogen as paired e- in oxygen in 2p subshell repel each other so less energy required to remove one of them

Question 8

define IE2

Answer 8

• the energy required to remove 1 mole of electrons from 1 mole of gaseous 1+ ions
• X^1+ —> X^2+ + e-
• endothermic only (more endothermic than 1IE) because energy needed to overcome the attraction between the positive ion and negative electron

Question 9

why is IE1 of Aluminium lower than that of Mg

Answer 9

e- removed in Al from more distant 3p (which is higher in energy) rather than 3s subshell in Mg

Question 10

why is IE1 of Sulfur lower than that of Phosphorus (3p)

Answer 10

less stable paired 3p orbital e- removed in sulfur

Question 11

what’s increase in nuclear charge outweighed by

Answer 11

increase in shielding

Question 12

what are giant covalent lattices

Answer 12

networks of atoms bonded by strong covalent bonds

Question 13

examples of giant covalent lattices

Answer 13

• diamond
• graphite
• graphene
• silicon

Question 14

why are giant covalent lattices insoluble, very hard and have high melting + boiling points

Answer 14

strong covalent bonds between atoms require lots of energy to overcome, so they’re not broken easily

Question 15

why are giant covalent lattices generally electrical insulators

Answer 15

no charged particles that are free to move

Question 16

why is diamond strong

Answer 16

• 4 strong covalent bonds per atom

Question 17

why is graphite slippery

Answer 17

• weak London forces between layers

Question 18

why can graphite and graphene conduct electricity

Answer 18

• conducts electricity as delocalised e- can move
• only 3 of the 4 available e- are in bonds so each C atom has one free e-

Question 19

graphite vs graphene

Answer 19

Graphene is simply one atomic layer of graphite

Question 20

which elements have a giant metallic structure in periods 2 and 3

Answer 20

• Li to Be
• Na to Al

Question 21

Why does the melting point of giant metallic structures increase in periods 2 and 3

Answer 21

• charge on metal increases
• no of delocalised e- in lattice structure increases
• strength of metallic bonding increases
• Li to Be
• Na to Al

Question 22

which elements have a giant covalent structure in periods 2 and 3

Answer 22

• B to C
• Si

Question 23

why does the melting point of giant covalent structures increase in periods 2 and 3

Answer 23

• lots of strong covalent bonds needing lots of energy to break to melt
• B to C
• Si

Question 24

which elements are simple molecules in periods 2 and 3

Answer 24

• N to Ne
• P to Ar

Question 25

Why is the melting point of simple molecules low in periods 2 and 3

Answer 25

only weak London forces are overcome on melting