Why do first ionisation energies increase

As you go down a group in the Periodic Table ionization energies generally fall. You have already seen evidence of this in the fact that the ionization energies in period 3 are all less than those in period 2. Taking Group 1 as a typical example:. Why is the sodium value less than that of lithium? There are 11 protons in a sodium atom but only 3 in a lithium atom, so the nuclear charge is much greater. You might have expected a much larger ionization energy in sodium, but offsetting the nuclear charge is a greater distance from the nucleus and more screening.

Lithium's outer electron is in the second level, and only has the 1s 2 electrons to screen it. The sodium's outer electron is in the third level, and is screened from the 11 protons in the nucleus by a total of 10 inner electrons. In other words, the effect of the extra protons is compensated for by the effect of the extra screening electrons. The only factor left is the extra distance between the outer electron and the nucleus in sodium's case.

That lowers the ionization energy. Similar explanations hold as you go down the rest of this group - or, indeed, any other group. Apart from zinc at the end, the other ionization energies are all much the same. All of these elements have an electronic structure [Ar]3d n 4s 2 or 4s 1 in the cases of chromium and copper. The electron being lost always comes from the 4s orbital. As you go from one atom to the next in the series, the number of protons in the nucleus increases, but so also does the number of 3d electrons.

The 3d electrons have some screening effect, and the extra proton and the extra 3d electron more or less cancel each other out as far as attraction from the center of the atom is concerned. In each case, the electron is coming from the same orbital, with identical screening, but the zinc has one extra proton in the nucleus and so the attraction is greater.

There will be a degree of repulsion between the paired up electrons in the 4s orbital, but in this case it obviously isn't enough to outweigh the effect of the extra proton. The lower the ionization energy, the more easily this change happens:. You can explain the increase in reactivity of the Group 1 metals Li, Na, K, Rb, Cs as you go down the group in terms of the fall in ionization energy.

Whatever these metals react with, they have to form positive ions in the process, and so the lower the ionization energy, the more easily those ions will form. The danger with this approach is that the formation of the positive ion is only one stage in a multi-step process.

For example, you would not be starting with gaseous atoms; nor would you end up with gaseous positive ions - you would end up with ions in a solid or in solution.

The energy changes in these processes also vary from element to element. Ideally you need to consider the whole picture and not just one small part of it. However, the ionization energies of the elements are going to be major contributing factors towards the activation energy of the reactions. Remember that activation energy is the minimum energy needed before a reaction will take place.

The lower the activation energy, the faster the reaction will be - irrespective of what the overall energy changes in the reaction are. The fall in ionization energy as you go down a group will lead to lower activation energies and therefore faster reactions. Jim Clark Chemguide. Trends of first ionization energies First ionization energy shows periodicity. Factors affecting the size of ionization energy ionization energy is a measure of the energy needed to pull a particular electron away from the attraction of 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.

The 3p electrons in phosphorus are all unpaired. In sulfur, two of the 3p electrons are paired. There is some repulsion between paired electrons in the same sub-shell, so the force of their attraction to the nucleus is reduced. This means that less energy is needed to remove one of these paired electrons than is needed to remove an unpaired electron from phosphorus.

All Rights Reserved. First ionisation energy across period 3. Learning outcomes After studying this page, you should be able to: describe and explain the trend in first ionisation energy across period 3. First ionisation energy The table shows first ionisation energy values for the elements Na to Ar. If you aren't reasonable happy about atomic orbitals and electronic structures you should follow these links before you go any further. Worried about moles?

Don't be! For now, just take it as a measure of a particular amount of a substance. It isn't worth worrying about at the moment. 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 which you would consider very low up to 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. 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. 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 more protons there are in the nucleus, the more positively charged the nucleus is, and the more strongly electrons are attracted to it. Attraction falls off very rapidly with distance. An electron close to the nucleus will be much more strongly attracted than one further away. Consider a sodium atom, with the electronic structure 2,8,1. There's no reason why you can't use this notation if it's useful!

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. This lessening of the pull of the nucleus by inner electrons is known as screening or shielding.

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.

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.

Hydrogen has an electronic structure of 1s 1. 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 kJ mol Helium has a structure 1s 2. 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 kJ mol -1 is much higher than hydrogen, because the nucleus now has 2 protons attracting the electrons instead of 1. Lithium is 1s 2 2s 1. 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 1s 2 electrons. Lithium's first ionisation energy drops to kJ mol -1 whereas hydrogen's is kJ mol Talking through the next 17 atoms one at a time would take ages.

We can do it much more neatly by explaining the main trends in these periods, and then accounting for the exceptions to these trends. The first thing to realise is that the patterns in the two periods are identical - the difference being that the ionisation energies in period 3 are all lower than those in period 2. In the whole of period 2, the outer electrons are in 2-level orbitals - 2s or 2p.

These are all the same sort of distances from the nucleus, and are screened by the same 1s 2 electrons.