Lesson 4 : Ionisation Energy

Defining first ionisation energy

Definition
The first ionisation energy is the energy required to remove the most loosely held electron from one mole of gaseous atoms to produce 1 mole of gaseous ions each with a charge of 1+.

This is more easily seen in symbol terms.
It is the energy needed to carry out this change per mole of X

Things to notice about the equation
The state symbols - (g) - are essential. When you are talking about ionisation energies, everything must be present in the gas state.

Ionisation energies are measured in kJ mol-1 (kilojoules per mole). They vary in size from 381 (which you would consider very low) up to 2370 (which is very high).

All elements have a first ionisation energy - even atoms which don't form positive ions in test tubes. The reason that helium (1st I.E. = 2370 kJ mol-1) doesn't normally form a positive ion is because of the huge amount of energy that would be needed to remove one of its electrons.
Patterns of first ionisation energies in the Periodic Table

The first 20 elements

First ionisation energy shows periodicity. That means that it varies in a repetitive way as you move through the Periodic Table. For example, look at the pattern from Li to Ne, and then compare it with the identical pattern from Na to Ar.

These variations in first ionisation energy can all be explained in terms of the structures of the atoms involved.

Factors affecting the size of ionisation energy

Ionisation energy is a measure of the energy needed to pull a particular electron away from the attraction of the nucleus. A high value of ionisation energy shows a high attraction between the electron and the nucleus.

The size of that attraction will be governed by:

The charge on the nucleus.
The more protons there are in the nucleus, the more positively charged the nucleus is, and the more strongly electrons are attracted to it.
The distance of the electron from the nucleus.
Attraction falls off very rapidly with distance. An electron close to the nucleus will be much more strongly attracted than one further away.

If the outer electron looks in towards the nucleus, it doesn't see the nucleus sharply. Between it and the nucleus there are the two layers of electrons in the first and second levels.

The 11 protons in the sodium's nucleus have their effect cut down by the 10 inner electrons.

The outer electron therefore only feels a net pull of approximately 1+ from the centre. This lessening of the pull of the nucleus by inner electrons is known as screening or shielding.

Warning! Electrons don't, of course, "look in" towards the nucleus - and they don't "see" anything either! But there's no reason why you can't imagine it in these terms if it helps you to visualise what's happening. Just don't use these terms in an exam! You may get an examiner who is upset by this sort of loose language.

Whether the electron is on its own in an orbital or paired with another electron.
Two electrons in the same orbital experience a bit of repulsion from each other. This offsets the attraction of the nucleus, so that paired electrons are removed rather more easily than you might expect.

Explaining the pattern in the first few elements

Hydrogen has an electronic structure of 1s1. It is a very small atom, and the single electron is close to the nucleus and therefore strongly attracted. There are no electrons screening it from the nucleus and so the ionisation energy is high (1310 kJ mol-1).
Helium has a structure 1s2. The electron is being removed from the same orbital as in hydrogen's case. It is close to the nucleus and unscreened. The value of the ionisation energy (2370 kJ mol-1) is much higher than hydrogen, because the nucleus now has 2 protons attracting the electrons instead of 1.

Lithium is 1s22s1. Its outer electron is in the second energy level, much more distant from the nucleus. You might argue that that would be offset by the additional proton in the nucleus, but the electron doesn't feel the full pull of the nucleus - it is screened by the 1s2 electrons.
You can think of the electron as feeling a net 1+ pull from the centre (3 protons offset by the two 1s2 electrons).

If you compare lithium with hydrogen (instead of with helium), the hydrogen's electron also feels a 1+ pull from the nucleus, but the distance is much greater with lithium. Lithium's first ionisation energy drops to 519 kJ mol-1 whereas hydrogen's is 1310 kJ mol-1.
The patterns in periods 2 and 3

*The general trend is for ionisation energies to increase across a period