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. 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. The major difference is the increasing number of protons in the nucleus as you go from lithium to neon. That causes greater attraction between the nucleus and the electrons and so increases the ionisation energies.
In fact the increasing nuclear charge also drags the outer electrons in closer to the nucleus. That increases ionisation energies still more as you go across the period. Note: Factors affecting atomic radius are covered on a separate page. Ionization energies increase across a Period, and decrease down a Group. Explanation: The first ionization energy is the energy required to produce a mole of gaseous ions and a mole of gaseous electrons from a mole of gaseous atoms.
Jun 12, Explanation: Two properties are important in determining ionization energies: i nuclear charge; and ii shielding by other electrons. Related questions How does ionization energy relate to reactivity? What is ionization energy measured in? What are the first and second ionization energies?
How does ionization energy increase? How does ionization energy change down a group? How do trends in atomic radius relate to ionization energy?
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