ionisation energy

Ionisation energy

The first ionisation energy is the amount of energy required to remove 1 mole of electrons from an isolated atom in the gaseous state. It can be represented by the equation:

X(g) → X+(g) + e

This process will obviously be an endothermic one, energy will have to be provided to remove the electron from the attractive force it feels from the positively charged nucleus. The first ionisation energies varies considerable for different elements. The three factors that you must consider when discussing ionisation energy are:

  1. the size of the nuclear charge, the larger the number of positively charged protons present in the nucleus then the greater will be the attraction for the electrons.
  2. The further away the electrons are from the nucleus then the easier they will be to remove since the force of attraction from the positvely charged protons in the nucleus will decrease with distance.
  3. The last factor to consider is shielding. The electrons in the valence shell (outer shell) will not feel the full affect of the positively charged nucleus because the inner or core electrons will effective shield or screen the nucleus from them. This shielding effect will reduce the size of the attractive force from the nucleus that the electrons feel and so it will require less energy to remove them.

Trends in the ionisation energy for the group2 elements

The first ionisation energies for the elements in group 2 of the periodic table are shown in the bar chart below. The general trend is fairly obvious, as we go down group 2 from the elements beryllium to barium the ionisation energy drops.

ionisation energies for the group 2 elements

To help explain this trend in the ionisation energies we need to consider the electronic configuration for the group 2 elements, these are shown in the table below:

element atomic number electron configuration
beryllium 4 1s22s2
magnesium 12 1s22s22p63s2
calcium 20 1s22s22p63s23p64s2
strontium 38 1s22s22p63s23p64s23d104p65s2
barium 56 1s22s22p63s23p64s23d104p65s24d105p66s2
As we move down group 2 from one element to the next a new electron shell is added and the electrons in the last shell or valence shell are in a higher principal energy level and so will be further from the nucleus. The size of the nuclear charge increases as we descend group 2 but the increasing nuclear charge is offset by the fact that the electrons in the valence ns sub-shell are further from the nucleus.

We can carry out a rough calculation to get an idea of the actual effective nuclear charge that the valence shell electrons will feel by simply subtracting the number of electrons in the lower electron shells from the nuclear charge e.g. Be has 4 protons, so its nuclear charge is 4+. It has 4 electrons in total, 1s22s2. There are 2 valence electrons in the 2s sub-shell, these being in the same shell will shield each other very weakly, so we will assume all the shielding comes from the inner 1s2 electrons. These 2 electrons will shield 2 protons. This means that the outer electrons will feel an effective nuclear charge of only 2+. We can carry out a similar calculation for all the group 2 elements, as shown in the table below:

element atomic number (nuclear charge) number of inner screening electrons number of valence electrons effective nuclear charge felt by valence electrons
beryllium 4 2 2 2+
magnesium 12 10 2 2+
calcium 20 18 2 2+
strontium 38 36 2 2+
barium 56 54 2 2+

This means that the valence electrons in all group 2 elements will feel an effective nuclear charge of 2+, but of course as we descend the group the distance from the nucleus to the valence electrons increases greatly, so much less energy will be required to separate the outer valence electrons when they are further from the positively charged nucleus, which means that the ionisation energy will get lower as the atoms in group get larger. We can show this simply as:

effective nuclear charge for group 2 metals

Successive ionisation energies

Ionisation energies are a good source of evidence for the presence of electron shells and sub-shells. So far we have only considered the enthalpy changes for the first ionisation energy of an element:

X(g) → X+(g) + e
However there is no reason to stop at removing just one electron, we can continue and remove more. The second ionisation energy of an element can be represented by the change:
X+(g) → X2+(g) + e
This is the enthalpy change (amount of heat energy required) to remove 1 mole of electrons from 1 mole of gaseous ions. The third ionisation energy would be:
X2+(g) → X3+(g) + e
As you might expect the ionisation energy required to remove successive electrons from an increasingly positively charged ions increases with each additional electron removed. As As an example consider the ionisation energies for aluminium, atomic number 13, with an electronic configuration: 1s22s22p63s231. The ionisation energy required to remove the first 7 electrons from aluminium is shown in the table below. As you can see the more electrons that are removed the more energy is required, however it is not a stepwise or steady increase.

ionisation energy 1st 2nd 3rd 4th 5th 6th 7th
energy required/kJmol-1 578 1820 2750 11 570 14 840 18 375 23 299

As a further example consider sodium, Na: 1s22s22p631 . Sodium has 1 valence electron, the electron in the 3s sub-shell. Once this electron is removed the ion formed will have a noble gas (np6) electron arrangement, in this case the noble gas will be neon. Removal of further electron will mean removing an electron from the second electron shell, the 2p electrons. These inner or core electrons are much closer to the nucleus and will be much more tightly held by the electrostatic attraction to the nucleus this coupled with the fact that we will be removing an electron from a samller positively charged sodium ion means more energy will be needed. The 1st ionisation energy of sodium is 496k/mol the second is 4560kJ/mol, quite an increase! This would be good evidence for the existence of electron shells within atoms. A similar thing is found with Mg, 1s2222p63s2. Magnesium has 2 valence electrons in the last shell, the 3s shell. You should be able to predict that removing the first two electrons, the valence electrons that would be lost in the normal chemical reactions of magnesium will require some energy but not too much. Once these two electrons are lost then magnesium will have a noble gas configuration (the same as Neon). However to remove a third electron would mean removing one from the second principal energy level, this will require a large increase in energy. The table below give the values for the first three ionisation energies of magnesium. This data provides clear evidence for electron shells. Here we have 2 electrons which are relatively easy to removed followed by a third which requires a huge increase in energy to remove:

ionisation energy 1st 2nd 3rd
energy required/kJmol-1 740 1819 7737

Key Points

Practice questions

Check your understanding - Questions on ionisation energies.