3.1.1 Periodicity

Question 1

How many sections is the periodic table divided into

What are they

Answer 1

4 (s,p,d+f block)

Question 2

Where is s block

Answer 2

Furthest to left and 2 upper sections

Question 3

Where is p block

Answer 3

Furthest to right

Question 4

Where is d block

Answer 4

Middle block

Question 5

Where is f block

Answer 5

Bottom block

Question 6

Where are the metals + non-metals around the stepped line

Answer 6

Metals are left of stepped line

Non-metals to the right of stepped line

Question 7

What are borderline elements classed as as they’ve got both metallic + non-metallic properties e.g silicon

Answer 7

Metalloids/semi-metals

Question 8

What’s a group

Answer 8

Vertical column of elements going down

Question 9

What’s a period

Answer 9

Horizontal row of elements going across

Question 10

What does periodic mean

Answer 10

Recurring regularly (a pattern)

Question 11

What are the 3 trends we look at

Answer 11

1st ionisation energy across a period
Atomic radius
Melting + boiling points across a period

Question 12

What 3 factors can explain al ionisation energy trends

Answer 12

Atomic radius (atom size)
Nuclear charge (proton number)
Shielding

Question 13

What does shielding take priority over

What does this mean

Answer 13

Shielding takes priority over increasing nuclear charge so down a group the ionisation energy increases as shielding increases, despite increased nuclear charge

Question 14

What’s the radius of an atom defined as

Answer 14

Half the distance between the centres of pairs of identical atoms, covalently bonded

Question 15

What happens to atomic radius across a period

Why

Answer 15

It decreases as protons are getting added but shielding stays the same so there’s a stronger attraction between electrons + nucleus

Question 16

What’s happens to atomic radius down a group

Why

Answer 16

It increases as the elements get 1 extra shell of electrons compared to the one before, so the outermost electrons are further from the nucleus

Question 17

Are atomic radius of metal atoms/ their corresponding ions smaller
Why

Answer 17

Ions

As they lose an electron, so attraction to nucleus is stronger

Question 18

Why are non-metal ions larger than their corresponding atoms

Answer 18

As they gain an electron , so attraction to nucleus slightly decreases

Question 19

What’s the only reason a boiling point is low

Answer 19

Only have to break intermolecular forces within simple covalent structures

Question 20

What have higher melting and boiling points (2)

Why

Answer 20

Metals
Form giant metallic lattices, with strong attraction between metal ions and delocalised electrons

Giant covalent (macromolecular) lattice
Have many strong covalent bonds that must be broken

Question 21

What is ionisation

Answer 21

Removing an electron from an atom, forming positive ion

Question 22

Define 1st ionisation energy

Eq

Answer 22

Energy needed to remove 1 mole of electrons from 1 mole of atoms in the gaseous state

M(g) -> M+(g) + 1e-

Question 23

Why are 1st + 2nd ionisation energies both endothermic

Answer 23

As energy must be put in to remove the negative electron from the nucleus’ attractive influence

Question 24

Define 2nd ionisation energy

Eq

Answer 24

Energy needed to remove 1 mole of electrons from 2 Mike of singly positively charged ions in the gaseous state

M+(g) -> M2-(g) + 1e-

Question 25

Why is the 2nd IE always greater than the 1st

Answer 25

As the electron is being removed from an already positively charged ion so is more strongly held

Question 26

How quickly are the electrons close to nucleus removed

Answer 26

Last (1s2)

Question 27

What does ionisation energy increase with

Why

Answer 27

Removal of each electron

As although charge stays same, remaining electrons feel charge stronger as electrons are being removed from a more positive species

Question 28

Why is there a big jump between 1st IE and 2nd

Answer 28

As it’s harder to remove the 2nd electron than the first - 2nd e- is closer to nucleus so has a stronger attraction due to less shielding

Question 29

Why do further out electrons feel less nucleus attraction

Answer 29

As there are more inner shells shielding them

Question 30

What does a big jump in energy mean

Answer 30

Moving to a shell closer to nucleus and electrons get harder to remove

Question 31

Why is electron pairing in an orbital mean electrons are easier to remove with less energy

Answer 31

As this creates repulsion between electrons

Question 32

Why does 1st IE decrease going down a group

Answer 32

Shielding increases so outer electrons feels a weaker nuclear attraction and needs less energy to remove

Question 33

3 main factors that affect ionisation energy

Answer 33

Atomic radius
Nuclear charge
Shielding

Question 34

What’s the ionisation energy like if the atomic radius is bigger

Answer 34

Lower IE as weaker attraction to nucleus

Question 35

How does nuclear charge affect ionisation energy

Answer 35

Stronger nuclear charge, stronger the attraction so greater ionisation energy

Question 36

How does shielding affect ionisation energy

Answer 36

Greater shielding, weaker attraction to nucleus so lower the ionisation energy

Question 37

What affects IE going across a group

Answer 37

Increasing nuclear charge