3.1.1 periodicity

how are elements arranged in the periodic table?

increasing atomic number

in periods showing repeating trends in physical and chemical properties (periodicity)

in groups having similar chemical properties

where are the s-, p- and d- blocks found on the periodic table?

s- groups 1 and 2

d- between groups 2 and 3

p- groups 13-18

first ionisation energy

The first ionisation energy is the energy required to remove one electron from each atom in one mole of gaseous atoms to form gaseous ions.

three factors affecting ionisation energy

Shielding- as the number of electrons between the outer electrons and the nucleus increases, there is less attraction towards the nuclear charge, since electrons are repelling each other (inner shell repelling outer shell).
Atomic radius- attraction between the nucleus and electrons decreases with distance
Nuclear charge- the more portions there are in the nucleus, the more positively charged it is, so there is a stronger attraction with the electrons.

why does ionisation energy decrease down a group?

atomic radius increases

more electron shielding

why does ionisation energy increase across a period?

increased nuclear charge due to number of protons increasing, making it harder to loose an electron

decreased atomic radius due to electrons being pulled in closer to the nucleus

decrease in the 1st ionisation energy between group 2 and 3

Group 3 elements have their outermost electron in a p-orbital

group 2 elements have theirs in an s-orbital

P-orbitals have a slightly higher energy level than s-orbitals and so are further away from the nucleus, making them easier to remove.

decrease in the first ionisation energy between group 5 and 6 elements

In group 5 each of the p-orbitals contains a single electron

In group 6 the outermost electron is now spin paired in the p-orbital

Electrons that are spin paired experiencing some repulsion- this makes the 1st outer electron easier to remove.

metallic bonding

The strong electrostatic attraction between cations and delocalised electrons

properties of diamond

each carbon covalently bonds with four others

tetrahedral arrangement

bond angles of 109.5

strong

high mp and bp (lots of energy needed to break covalent bonds)

does not conduct electricity

insoluble

properties of graphite

each carbon covalently bonded to three others

layers held together by weak london forces

delocalised electron means electricity can be conducted

properties of graphene

single layer of graphite

better at conducting electricity

properties of silicon?

brittle

semi conductor

high mp and no

relative mp and bp of a giant metallic lattice

Generally high mp and bp

Depends on the strength of the metallic bonds holding together the atoms in the lattice

lots of energy is needed to overcome the strong electrostatic attraction between cations and electrons

solubility of giant metallic lattices

They do not dissolve- any interactions would lead to a reaction

electrical conductivity of giant metallic lattices

They can conduct electricity in both solid and liquid states

When a voltage is applied across a metal, the delocalised electrons can move through the structure, carrying charge

relative mp and bp of a giant covalent lattice

high mp and bp due to strong covalent bonds

high temperatures are necessary to provide the large quantity of energy needed to break the strong covalent bonds

solubility of a giant covalent lattice

insoluble, due to strong covalent bonds, which cannot be broken by interaction with solvents

electrical conductivity of a giant covalent lattice

generally, non-conductors (except graphite, graphene)

In carbon (diamond) and silicon, all 4 outer shell electrons are involved in covalent bonds, so none are available for conducting electricity

In graphite and graphene, only 3 outer shell electrons are covalently bonded, leaving free electrons to conduct electricity.

periodic trend in mps from period 2-3

mp increased from group 1-4

sharp decrease between group 4-5

mp are low from 5-0

sharp decrease in ml marks a change from giant to simple molecular structures

on melting giant structures have sting forces to overcome so have high melting points

simple molecular structures have weak forces to overcome, so have much lower bps

this trend is repeated across period 3 downwards