How can electronegativity be used to distinguish between

Pauling derived the first electronegativity values by comparing the amounts of energy required to break different types of bonds. He chose an arbitrary relative scale ranging from 0 to 4. When the difference is very small or zero, the bond is covalent and nonpolar.

When it is large, the bond is polar covalent or ionic. The absolute values of the electronegativity differences between the atoms in the bonds H—H, H—Cl, and Na—Cl are 0 nonpolar , 0.

The degree to which electrons are shared between atoms varies from completely equal pure covalent bonding to not at all ionic bonding. This table is just a general guide, however, with many exceptions. For example, the H and F atoms in HF have an electronegativity difference of 1. Likewise, the Na and Cl atoms in NaCl have an electronegativity difference of 2.

The best guide to the covalent or ionic character of a bond is to consider the types of atoms involved and their relative positions in the periodic table. Bonds between two nonmetals are generally covalent; bonding between a metal and a nonmetal is often ionic. Some compounds contain both covalent and ionic bonds. However, these polyatomic ions form ionic compounds by combining with ions of opposite charge.

Bond polarities play an important role in determining the structure of proteins. The polarity of these bonds increases as the absolute value of the electronegativity difference increases.

So I go back down here, and I go ahead and put in a 1. And so that's a difference in electronegativity of 1. So we could consider this to be a polar covalent bond. This time, carbon is more electronegative than lithium.

So the electrons in red are going to move closer to the carbon atom. And so the carbon is going to have a little bit more electron density than usual. So it's going to be partially negative. And the lithium is losing electron density, so we're going to say that lithium is partially positive.

Now here, I'm treating this bond as a polar covalent bond. But you'll see in a few minutes that we could also consider this to be an ionic bond. And that just depends on what electronegativity values you're dealing with, what type of chemical reaction that you're working with. So we could consider this to be an ionic bond. Let's go ahead and do an example of a compound that we know for sure is ionic. Sodium chloride, of course, would be the famous example.

So to start with, I'm going to pretend like there's a covalent bond between the sodium and the chlorine. So I'm going to say there's a covalent bond to start with.

And we'll put in our electrons. And we know that this bond consists of two electrons, like that. Let's look at the differences in electronegativity between sodium and chlorine.

So I'm going to go back up here. I'm going to find sodium, which has a value of 0. So sodium's value is 0. Chlorine's is 3. That's a large difference in electronegativity. That's a difference of 2. And so chlorine is much more electronegative than sodium. And it turns out, it's so much more electronegative that it's no longer going to share electrons with sodium. It's going to steal those electrons.

So when I redraw it here, I'm going to show chlorine being surrounded by eight electrons. So these two electrons in red-- let me go ahead and show them-- these two electrons in red here between the sodium and the chlorine, since chlorine is so much more electronegative, it's going to attract those two electrons in red so strongly that it completely steals them. So those two electrons in red are going to be stolen by the chlorine, like that. And so the sodium is left over here.

And so chlorine has an extra electron, which gives it a negative 1 formal charge. So we're no longer talking about partial charges here. Chlorine gets a full negative 1 formal charge. Sodium lost an electron, so it ends up with a positive formal charge, like that. And so we know this is an ionic bond between these two ions. So this represents an ionic bond. So the difference in electronegativity is somewhere between 1.

So most textbooks we'll see approximately somewhere around 1. So if you're higher than 1. Lower than 1. But that doesn't always have to be the case. So we'll come back now to the example between carbon and lithium. So if we go back up here to carbon and lithium, here we treat it like a polar covalent bond.

But sometimes you might want to treat the bond in red as being an ionic bond. So let's go ahead and draw a picture of carbon and lithium where we're treating it as an ionic bond.

So if carbon is more electronegative than lithium, carbon's going to steal the two electrons in red. So I'll go ahead and show the electrons in red have now moved on to the carbon atom.

So it's no longer sharing it with the lithium. Carbon has stolen those electrons. And lithium is over here. So lithium lost one of its electrons, giving it a plus 1 formal charge. Carbon gained an electron, giving it a negative 1 formal charge. And so here, we're treating it like an ionic bond. The degree to which a given bond is ionic or covalent is determined by calculating the difference in electronegativity between the two atoms involved in the bond.

As an example, consider the bond that occurs between an atom of potassium and an atom of fluorine. Using the table, the difference in electronegativity is equal to 4. Since the difference in electronegativity is relatively large, the bond between the two atoms is ionic. Since the fluorine atom has a much larger attraction for electrons than the potassium atom does, the valence electron from the potassium atom is completely transferred to the fluorine atom.

The diagram below shows how difference in electronegativity relates to the ionic or covalent character of a chemical bond. A bond in which the electronegativity difference is less than 1. However, at this point we need to distinguish between two general types of covalent bonds.

A nonpolar covalent bond is a covalent bond in which the bonding electrons are shared equally between the two atoms. In a nonpolar covalent bond, the distribution of electrical charge is balanced between the two atoms. Figure 3. A nonpolar covalent bond is one in which the distribution of electron density between the two atoms is equal.

The two chlorine atoms share the pair of electrons in the single covalent bond equally, and the electron density surrounding the Cl 2 molecule is symmetrical. An example would be a bond between chlorine and bromine. Figure 4. In the polar covalent bond of HF, the electron density is unevenly distributed. There is a higher density red near the fluorine atom, and a lower density blue near the hydrogen atom.