 The last trend we're going to look at is electronegativity. Electronegativity is a measure of how strongly an atom will pull electrons towards itself when it forms a bond. Each element has a different electronegativity. While we can talk about the electronegativity of an individual type of atom, this property really only manifests when an atom is bonded to other atoms. It cannot be measured directly, but it's calculated from other atomic and molecular properties. A quick refresher on bonding. When a covalent bond forms between two atoms, two bonding electrons are shared between them. Just as in an individual atom there is an electrostatic attraction between the negative electrons and the positive nucleus, so in a covalent bond between two atoms, the bonding electrons are electrostatically attracted to the nuclei of both the atoms in the bond. However, the more electronegative atom attracts the bonding electrons more strongly, meaning it takes the lion's share of the electron density. This hydrogen fluoride molecule has two bonding electrons that hold the atoms together. The fluorine atom has significantly higher electronegativity than the hydrogen, and that means that those two bonding electrons spend more time close to the fluorine atom than they do to the hydrogen. The effect of this behaviour is that the fluorine atom effectively has a partial negative charge. The electrons spend more time around it, and the hydrogen, which is a bit electron deficient, has a partial positive charge. We represent this using a small delta negative, meaning less than a full negative charge, and small delta positive. This kind of bond is known as a polar covalent bond because there are shared electrons, but the electron density is unequally distributed between the two atoms. We'll talk more about this in a later video on bonding. So what are the trends in electronegativity? This periodic table shows higher electronegativities in a darker red colour. You may notice a couple of things. First electronegativity increases from left to right across the periods, and increases from bottom to top of each group, much the same trend as ionisation energy. In fact, fluorine is the element with the highest electronegativity. If you remember this, it's easy to remember the trend. In general, electronegativity increases in the direction of fluorine. You may also notice that this table has no values for some of the elements. Remember that unlike ionisation energy, electronegativity only makes sense in the context of a bond between two atoms. The element has to have bonded with another atom for its electronegativity to be calculated. Helium, neon and argon, because they're such stable atoms, have never been found to form bonds with other atoms, and the radioactive elements in period 7 are either too short-lived or have not been studied in this regard. So can we rationalise these trends? Well, yes, we can, and it may not come as a surprise that it comes down to the same factors that we discussed before in the context of atomic radius and ionisation energy. In fact, the situation is very similar to ionisation energy. With the first ionisation energy, we're essentially measuring the attraction between the nucleus of an atom and its outermost electron. With electronegativity, we're measuring the attraction between the nucleus of a bonded atom, that's an atom that's part of a molecule, and its bonding electrons. The situation is similar and so the same factors apply. As before, remember that you can't take these factors in isolation. All three have to be considered to explain the properties of an atom. As always, the nuclear charge is important. The more positively charged the nucleus, the stronger the attraction for bonding electrons. So more protons in the nucleus tends to mean that an atom has higher electronegativity. Second, the distance from the nucleus. When large atoms form bonds, the bonding electrons that sit between them are far from the nucleus. They occupy a similar region in space to the valence shell. So the electrostatic attraction between the nucleus and the bonding electrons is decreased. The greater the number of electron shells, or you could say the larger the atom, the lower the electronegativity. Third is the shielding effect. More electron shells means more shielding of the nuclear charge, so a lower attraction between nucleus and bonding electrons, and hence a lower electronegativity. So what's going on is that as you go across a period, the nuclear charge increases while the atom gets smaller, which means that the distance of the bonding electrons from the nucleus decreases. Remember that trend for atomic radius. Also, the shielding doesn't change that much, because you're staying in the same electron level. The total effect is that electronegativity increases. Going down a group, the nuclear charge still increases, but it's more than cancelled out by the increased radius of the atom, and the extra electron levels that provide shielding, so electronegativity decreases.