3.1.1 Periodicity

Question 1

Classification of elements into orbital blocks

Answer 1

the last electron in the final orbital decides where in the periodic table the element falls, which orbital it is in decides its placement in the table

Question 2

what is periodicity

Answer 2

repeating pattern across different periods

Question 3

atomic radius across a period

Answer 3

decrease
increased protons means increased nuclear charge
increased nuclear attraction for electrons in the same shell
shielding stays the same

Question 4

definition of first ionisation energy

Answer 4

energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions

Question 5

equation to represent 1st ionisation energy of hydrogen

Answer 5

H(g) -> H+(g) + e-

Question 6

factors that affect ionisation energy

Answer 6

1. nuclear attraction
2. atomic radius
3. electron shielding

Question 7

when does nuclear attraction increase

Answer 7

more protons in nucleus the greater the attraction

Question 8

when does atomic radius increase

Answer 8

larger the atom the further away the outer electrons are from the nucleus and the weaker attraction to the nucleus

Question 9

shielding of electrons

Answer 9

electron in outer shell is repelled by electrons in complete inner shell, weakening the attraction to the nucleus

Question 10

general pattern of successive ionisation energies

Answer 10

increases as more electrons are removed

Question 11

why does each successive ionisation energy get larger

Answer 11

the ion formed is smaller than the atom and the proton to electron ratio in the 2+ ion is greater than the 1+ ion
therefore the nuclear attraction is stronger

Question 12

how are ionisation energies linked to electronic structure?

Answer 12

if there is a big jump between 2 values, can show evidence for sub shells.
big jump as the shell is then closer to the nucleus and attracted much more strongly

Question 13

how can you tell which group an element is in when given successive ionisation energies

Answer 13

look for a big jump between two consecutive values
e.g. between 2 and 3
meaning the element must be in group 2 because the 3rd electron to be removed must be in a shell closer to the nucleus with less shielding so will have a larger ionisation energy

Question 14

general trend of first ionisation energies of the elements

Answer 14

generally increases across the group

Question 15

actual trend of first ionisation energies of the elements

Answer 15

increases

exception at Mg to Al and P to S

Question 16

explain exceptions of first ionisation energies of the elements

Answer 16

Mg to Al: Al is filling 3p subshell but Mg has electrons in 3s which are more difficult to remove. Electrons in 3p are higher in energy and shielded by the 3s electrons, only one electron in the orbital so isn’t paired
P to S: sulfur has a 4th electron in the p orbital, the electron sare starting to pair with opposite spin so there is increased repulsion between the electrons and so is easier to remove

Question 17

why does helium have the highest first ionisation energy

Answer 17

its first electron is in the first shell closest to the nucleus so has no shielding electrons, higher than hydrogen as there is one more proton

Question 18

trend in ionisation energies down a group

Answer 18

decreases

Question 19

why do first ionisation energies decrease down a group

Answer 19

as you go down the group the electrons are found in shells further away from the nucleus, increased shielding but the nuclear attraction decreases making it easier to remove 1 mole of electrons

Question 20

why do first ionisation energies increase across a period

Answer 20

electrons added to the same shell 
shielding stays the same 
number of protons increases
nuclear charge increases
distance from nucleus stays the same 
nuclear attraction increases
more difficult to remove electrons

Question 21

why does sodium have a much lower first ionisation energy than neon

Answer 21

sodium has its outer shell electron in a 3s subshell further from the nucleus and is more shielded, outer electron is easier to remove

Question 22

definition of metallic bonding

Answer 22

electrostatic attraction between positively charged ions and the sea of delocalised electrons

Question 23

3 main factors that determine metallic bond strength

Answer 23

1. Number of protons (more protons, greater charge, stronger the bond)
2. Charge of metal ion/delocalised electrons (more delocalised electrons the stronger the bond)
3. Size of the ion (smaller the ion the stronger the bond)

Question 24

which metallic bond is stronger Mg or Na?

Answer 24

Mg
more electrons in outer shell that are delocalised
charge on the ion is greater
Mg ion is smaller and has one more proton
stronger force of electrostatic attraction between the positive ion and the electrons so increased energy is required to break the bond
Mg will have a higher melting point

Question 25

2 examples of macromolecular compounds

Answer 25

diamond

graphite

Question 26

structure of diamond

Answer 26

macromolecular
tetrahedral arrangement of carbon atoms
4 covalent bonds per carbon atom

Question 27

structure of graphite

Answer 27

macromolecular
planar arrangement of carbon atoms
each carbon is bonded to 3 other carbons
1 electron per carbon is donated into the sea of delocalised electrons
which are located between the layers

Question 28

melting and boiling points of diamond and graphite

Answer 28

very high
due to strong covalent bonds
require high amounts of energy to break them

Question 29

boiling and melting points of macromolecular

Answer 29

high- because of many strong
covalent bonds
lot of energy to break the many
strong bonds

Question 30

boiling and melting points of metallic

Answer 30

high- strong electrostatic forces between positive

ions and sea of delocalised electrons

Question 31

macromolecular solubility in water

Answer 31

insoluble

Question 32

metallic Solubility in water

Answer 32

insoluble

Question 33

conductivity when solid macromolecular

Answer 33

diamond and sand: poor, because
electrons can’t move (localised)
graphite: good as free delocalised
electrons between layers

Question 34

conductivity when solid metallic

Answer 34

good: delocalised electrons can move through

structure

Question 35

conductivity when molten macromolecular

Answer 35

poor

Question 36

conductivity when molten metallic

Answer 36

good

Question 37

general description of giant metallic

Answer 37

shiny metal
Malleable as the positive ions in the lattice are all
identical. So the planes of ions can slide easily over
one another
-attractive forces in the lattice are the same
whichever ions are adjacent

Question 38

general trend in melting and boiling points across a period

Answer 38

increase until Si then decrease

Question 39

describe the trend in melting and boiling points across a period

Answer 39

increases Na to Al
peaks at Si
decreases to P
increases at S 
decreases Cl to Ar

Question 40

explain the trend in melting and boiling points across a period

Answer 40

Na-Al: high due to metallic bonding, high bond strength, increased energy required to break the bonds
Si: macromolecular, strong covalent bonds, more energy to break
P-Ar: simple molecular so require less energy to break the bonds

Question 41

melting and boiling points simple molecular

Answer 41

S8
P4
Cl2
Ar

decreases due to decrease in surface area, less london forces, less energy required to break the bonds