Ionisation energies

How to measure atomic radius

No way to measure it directly as electron clouds don’t have a clear cut off point. Can roughly measure it by looking at identical atoms that have formed pairs and take the atomic radius as half the distance between the nuclei of the two atoms

Trends in atomic radius across periods

Atomic radius decreases because there is an increase in proton number meaning there is an increase in nuclear charge so there is a bigger attraction between the nucleus and electrons pulling the electrons closer to the nucleus making the atomic radius smaller. The screening effect stays the same across the period

Trends in atomic radius down groups

Atomic radius increases down groups because the amount of shells the elements have increases meaning there is more screening between the nucleus and electron so there is less attraction between the nucleus and electrons meaning there’s less pull between them

First ionisation energy

The energy needed to remove one electron from each atom in one mole of a gaseous atoms to form one mole of a gaseous +1 ion

Ionic equation for first ionisation energy

X (g) —> X+ (g) + e-

First ionisation energy trend across a period

Tends to increase as the number of protons increases resulting in a bigger nuclear charge so more attraction between the nucleus and outer electron also resulting in a decrease in atomic radius so more energy is required to remove the first outermost electron. Screening affect stays the same across the period

Drop from Mg to Al in first ionisation energy

Al has its outermost electron in a 3p sub shell whereas Mg has its outermost electron in a 3s sub shell. It’s easier to remove an electron from a 3p sub shell than a 3s sub shell so less energy is required to remove Al’s outermost electron

Drop from Ph to S in first ionisation energy

Phosphorus has one electron in its 3s orbital whereas sulphur has 2 in its 3s orbital and these two electrons repel each other so there is more screening so less energy is required to remove sulphurs outermost electron

First ionisation energy trends down a group

Tends to decrease down a group as there is an increase in atomic radius as well as an increase in the amount of sub shell each element has so an increase in screening this cancels out the increase in nuclear charge resulting in less energy needed to be used to remove the outermost electrons

Trends in successive IE’s within elements

There is a general increase at first because as each electron is removed the ion becomes increasingly positive and the atomic radius decreases the shielding effect decreases and the attraction between the electrons and nucleus increases making it harder to loose and electron

Example successive IE ionic equation

X+ (g) —> X2+ (g) + e- ( the resulting ion tells you the successive IE it is)

How to identify element from successive IE values

Work out the number of electrons in the outermost shell as these have a gradual increase in IE values which tells you what group the element is in