Electronegativity

Covalent bonding

A covalent bond is usually formed when 2 non-metals atoms join and share a pair of electrons EQUALLY between them. You should be familiar with covalent bonds from your gcse chemistry course. The diagram below shows a covalent bond formed between 2 atoms of hydrogen. You should recall that the two electrons in the covalent bond between the hydrogen atoms are attracted to each of the positively charged hydrogen nuclei. The two negatively charged electrons will feel the same attractive force from each nucleus since both nuclei contain 1 positively charged proton. This means that that the two electrons are shared equally between the two hydrogen atoms. You will no doubt have learned from your gcse course that a covalent bond involves the equal sharing of a pair of electrons. This is outlined in the diagram below:

However what happens if the molecule contains different atoms? Will the electrons in the covalent bond still be shared equally this time? Consider as an example a molecule of hydrogen fluoride.

The fluorine atom has a much larger nuclear charge (+9) compared the (+1) nuclear charge in the hydrogen nucleus, so the fluorine atom will be better at attracting the electrons in the covalent bond than hydrogen atom ; we say that fluorine atom is more electronegative than the hydrogen atom. Electronegativity is the ability of an atom in a covalent bond to attract electron density towards it.

Factors affecting electronegativity

The Electronegativity of an element depends on a number of factors. One factor is the effective nuclear charge that the electrons in the valency shell feel, since these are the electrons involved in actually forming covalent bonds; for example fluorine has an electron arrangement of 2,7 or 1s²2s²2p⁵. The 2 inner shell electrons in the 1s sub-shell will screen or shield 2 protons from the outer shell electrons and since electrons in the same shell are not particularly effective at shielding each other this means that the outer valency shell electrons will feel an effective nuclear charge of 9-2=+7. For hydrogen the effective nuclear charge is only +1 since it contains only the 1 proton and 1 electron.

As we cross any period in the periodic table the number of inner shell electrons stays the same but the nuclear charge is increasing. Since the nuclear charge is increasing and no additional energy levels are being added the size of the atoms will decrease as the extra protons in the nucleus pull the electrons closer to it.

This means that the electronegativity will increase across a period in the periodic table. This also means that the elements in the right side of the period will be better at attracting electrons in a covalent bond towards them- that is they have higher electronegativity values.

The size of the atom will also have an effect on its electronegativity value. As we descend a group in the periodic table the effective nuclear charge remains the same but the atoms obviously get larger; this means that the nucleus is further away from any electrons in a covalent bond and so will be less able to attract any electrons in the covalent bond electron density. This means that the electronegativity decreases down a group e.g. consider the two alkali metals lithium and potassium.

Potassium atoms (₁₉K) are obviously larger than lithium atoms (₃Li) and also have a larger nuclear charge (+19) however there are 18 electrons between the nucleus and the outer 4s¹ electron. These 18 electrons will effectively shield or screen out 18 protons or 18 positive charges. This means that the outer 4s¹ electrons will only feel an effective nuclear charge of +1 and NOT +19.

A similar argument can be made for lithium; the inner 1s² electrons will shield or "cancel" out the charge from 2 protons in the nucleus; so as with potassium the outer 2s¹ electron in lithium will feel an effective nuclear charge of +1. However the outer electron in lithium is closer to the nuclear charge so it will feel a greater attraction from the effective +1 nuclear charge. The outer valency electron in potassium is much further from the nucleus so the attraction to the +1 effective nuclear charge this time will be less, this is shown in the diagram above and summarised in the table below.

The trends in the electronegativity values in a group and across a period in the periodic table as shown in the image below:

Small atoms with large effective nuclear charges have large electronegativity values. The most electronegative element in the periodic table is fluorine while the least electronegative element is francium. The image above illustrates the trends in electronegativity across the periodic table. As we descend a group the atomic radius of the atom increases while the effective nuclear charge remains constant; this means that electronegativity decreases down a group in the periodic table. As we cross a period in the periodic table the atomic radius decreases while the effective nuclear charge increases; this means that the electronegativity increases across a period. The alkali metals in group 1 of the periodic table have low electronegativity values whereas the halogens in group 7 have higher electronegativity values.

Electron affinity, ionisation energy and electronegativity

Electron affinity is the ability of an isolated atom to attract an electron and the ionisation energy is a measure of how easy or hard it is to remove an electron from an atom. So one of the ways to measure electronegativity is to take an average of the electron affinity and ionisation energy. A scale can be produced from this and a number assigned to each element to represent its electronegativity (note the scale or number has no units). On this scale the most electronegative element fluorine is assigned a value of 4 and all other elements electronegativity values are measured relative to this. The table below lists some values of electronegativity for the elements you are likely to meet. You can see that N, O and F are the most electronegative elements in the periodic table. It is also clear to see the trends in the electronegativity values across a period and down a group in the periodic table.

It is worth mentioning that if you look online or in textbooks you may see slightly different values of electronegativity for some elements; this is simply due to the way in which it is measured. However the same patterns or trends should be obvious no matter which method is chosen.

There are no values listed for the noble gases; this is simply because they rarely form covalent compounds and have no affinity for electrons.

Key Points

• Electronegativity is the ability or power of an atom to attract or pull electron density in a covalent bond towards it.
• Electronegativity increases across a period in the periodic table because the nuclear charge is increasing and the size of the atoms is decreasing. No additional shielding takes place since the addition of extra electrons as we cross a period go into the same principal energy level and electrons within the same principle energy level are poor at shielding each other from the nuclear charge.