7 Periodicity

what do groups show and how does this affect their chemical properties

• same number of e- in outer shell

• same number of e- in sub shells

• similar chemical properties

what do periods show

• number of electron shells

periodicity def

the trends seen within groups and periods or a ‘repeating periodic pattern’

e.g. electron configuration, ionisation energy, structure, mpts

what factors affect atomic radii

atomic number

• distance between nucleus and outer electrons

• number of shells

how does atomic radius change across periods and groups

• atomic radius decreases across a period

• atomic radius increases down a group

what do blocks (s-block, p-block) show on a period table

the highest energy subshell that is filled

what happens to atomic radius across a period and why

• atomic radius decreases across a period

• nuclear charge increases while outer electrons experience the same shielding

• greater nuclear attraction on the outer shell electrons

metallic bonding

strong electrostatic attraction between positive metal cations and a sea of delocalised electrons

metallic bonding diagram

what properties are looked at in this chapter 😭😭

• electron configuration

• ionisation energy

• structure

• melting points

trend of electron configuration across a period

each period starts w an electron in a new highest energy shell

• across period 2, the 2s sub-shell fills w 2e-, followed by the 2p sub-shell w 6e-

• across period 3, same pattern is repeated

• across period 4 the highest shell number is n = 4, from the n = 4 shell only the 4s and 4p sub-shells r occupied as the 3d sub-shell is involved

trend of electron configuration down a grp

• elements in each group have atoms w the same no. e- in their outer shell

• elements in each grp also have atoms w the same no. e- in each sub0shell

• this gives elements in the same grp similar chemistry

how are elements divided into blocks

corresponding to their highest energy sub-shell

blocks on periodic table pic

first ionisation energy def

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

factors that affect 1st ionisation energy

• atomic radius

• electron shielding

• nuclear charge

(mention all in exam Qs)

how does atomic radius affect IE

• if larger atomic radius

• outer electron is further from nucleus

• e- has less nuclear attraction

• less energy needed to remove

• easier to remove outer e-

how does e- shielding affect IE

• shielding effect means more e-

• therefore more repulsion

• outer e- has less nuclear attraction

• less energy required to remove it

• easier to remove outer e-

how does nuclear charge affect IE

• more protons → higher nuclear charge

• outer e- has a greater nuclear attraction

• more energy required to remove the outer e-

• harder to remove outer e-

1st IE equation (using X as a random element)

2nd IE def

energy required to remove 1 e- from each ion in 1 mole of gaseous 1+ ions of an element to form 1 mole of gaseous 2+ ions

why does Al have a higher bpt than Na

• Al 3+ ions have a greater charge than Na+ ions

• therefore more delocalised electrons

• stronger electrostatic attractions between cation and delocalised electrons

• more energy needed to overcome

• higher bpt

structure of Si

giant covalent structure

• breaking strong covalent bonds

• lots of energy to overcome

how many other Si is Si covalently bonded to

each Si is covalently bonded to 4 others

how does P travel

P₄

how does S travel

S₈

why does P and S have a higher bpt than Cl and Ar all in period 3

P₄

• 4 atoms

• more e-

• stronger LF

• more energy required to overcome

S₈

• much higher mpt&bpt expected of a simple molecular structure

• 8 atoms

• many more e-

• stronger LF

• more energy required to overcome

what type of IMF r present from P₄ to Ar

• no polar bonds therefore no permanent dipole-dipole interactions

• no HB

• only LF

how does bpt change down grp17 (7) (halogens)

• all simple covalent/molecular

• all diatomic

• all LF

• down the grp atomic radius increases

• more e-

• stronger LF

• more energy to overcome

• bpt increases down grp7

bpt down grp1

• bpt decreases down group

• metallic lattice structure

• all 1+ ions attracting 1e-

• delocalised e- have more attraction higher up the grp

• closer to nucleus due to smaller atomic radius

• stronger electrostatic attraction

• more energy to overcome (higher up the grp)

atomic radii across period

• decreases across a period

atomic radius exam q ans

• atomic radius decreases across period

• nuclear charge increases while outer electrons experience the same shielding

• greater nuclear attraction on the outer shell electrons

1st IE down grp1 graph

• decreases

• easier to remove outer e- down grp1

explain why 1st IE decreases down a group (1)

1. first IE decreases down a group where the outer e- is easier to remove

2. although nuclear charge increases, so does the atomic radius

3. outer electron further from nucleus

4. less nuclear attraction

5. more electron shielding therefore greater shielding effect therefore greater repulsion between electrons

6. less nuclear attraction to outer electron

7. less energy require to remove the outer electron

1st IE across period 2 and 3

explain 1st IE trend across a period

1. across a period the 1st IE increases - harder to remove the outer e-

2. across a period nuclear charge increases - more nuclear attraction

3. similar electron shielding because the outer electrons are in the same energy level

4. atomic radius decreases across a period

5. overall the outer e- has more attraction to the nucleus across a period - more energy required to remove it

where are there drops in period 2 and why

grp2&3 and group5&6, Be → B and N → O

• B (grp3) outer e- is in a higher energy level than Be (grp2)

• more e- shielding

• less nuclear attraction

• less energy required to remove despite higher nuclear charge

explain y there is a drop from N → O

• for O (grp6) one p-orbital is doubly filled

• repulsion between 2e- in orbital

• outer e- has less nuclear attraction

• despite higher charge

• less energy to remove

why do successive IEs increase within a shell

• bcuz u r removing an e- from an increasingly positive ion

• greater nuclear attraction

electron configuration for chromium

1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵ 4s¹

electron configuration for copper

1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d¹⁰

1st IE of oxygen

O_(g) -> O_(g)⁺ + e^-

describe the structure of metals

billions of metal atoms held together by metallic bonding in a giant metallic lattice

• in a solid metal structure each atom has donated its negative outer shell electrons to a shared pool of electrons which are delocalised throughout the whole structure

• the delocalised e- are mobile and able to move throughout the structure.

why do metals have high mpts and bpts

• most metals high temperatures are needed to provide the large amount of energy needed to overcome the strong electrostatic attraction between the cations and electrons

• strong attractions → high mpts and bpts

why can metals conduct electricity when solid or molten

• when a voltage is applied across a metal the delocalised e- can move through the structure carrying charge

• has mobile charge carriers

are metals soluble

no, any interactions w the charges in a metallic lattice would lead to a reaction rather than dissolving

why do giant covalent structures have high mpts and bpts

• has strong covalent bonds

• high temperatures needed to provide the large amount of energy needed to break the strong covalent bonds

would diamond be soluble in water and why

• no

• giant covalent lattices are insoluble in almost all solvents

• the covalent bonds holding together the atoms in the lattice are too strong to be broken by interactions w solvents

why can’t diamond conduct electricity

all 4 outer shell e- are involved in covalent bonding

none available for conducting electricity

what would a dot and cross diagram of structure and bonding in carbon (diamond) look like

how many carbons is each carbon bonded to in graphite, what is the bond angle ?

3, 120°

how can graphite conduct electricity if its a non-metal

• the bonding in the hexagonal layers only uses 3 of carbon’s 4 outer shell e-

• the spare e- is delocalised between layers

• therefore there is a mobile charge carrier

• can conduct electricity

what is graphene

a single layer of graphite

can graphene conduct electricity

yes, there is 1 delocalised e- that’s free to move

structure of graphite

what is between the hexagonal layers of graphite

weak forces

dot and cross of a graphene layer

why is there a decrease in mpt from group4-5

• the sharp decrease in mpt is due to the change from a giant covalent structure to simple molecular structure

• giant covalent structures have strong covalent bonds between atoms, higher temps needed to provide the energy to break them

• simple molecular structures only have weak LF between molecules that don’t take much energy to separate