The dipole moment is a chemical property that indicates how heterogeneously electric charges are distributed in a molecule. It is expressed in Debye units, 3.33 · 10-30 C · m, and generally its values range from 0 to 11 D.
Highly polar compounds tend to have large dipole moments; while the apolar ones, small dipole moments. The more polarized the electrical charges are in a molecule, the greater its dipole moment; that is, there must be a region rich in electrons, δ-, and another poor in electrons, δ+.
The dipole moment, μ, is a vector magnitude, so it is affected by the angles of the bonds and, in general, by the molecular structure.
When the molecule is linear, it can be compared to a two-color eraser. Its negative end δ-, would correspond to the color red; while the positive, δ +, would be the blue color. As the magnitudes of the negative charges at the δ- pole increase, and the distance that separates it from δ +, the dipole moment increases.
Chemically, this means that the greater the difference in electronegativity between two atoms, and the longer the distance that separates them, the greater the dipole moment between them..
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It is considered a covalent bond between two atoms, A and B:
A-B
The distance between the positive and negative partial charges is already defined by the length of their bond:
TOδ+-Bδ-
Because protons and electrons have the same magnitude of electric charge but with opposite signs, 1.6 · 10-19C, this is what is taken into account when evaluating the dipole moment between A and B using the following equation:
μ = δd
Where μ is the dipole moment, δ is the charge of the electron without the negative sign, and d the length of the link expressed in meters. For example, assuming that d has a value of 2 Å (1 10-10m) the dipole moment, μA-B will be:
μA-B = (1.6 10-19C) · (2 · 10-10m)
= 3.210-29C m
But since this value is very small, the Debye unit is used:
μ = (3,210-29C · m) · (1 D / 3.33 · 10-30 C m)
= 9.60 D
This value of μA-B could give to suppose that the bond A-B is more ionic than covalent.
To calculate the dipole moment of a molecule, all the dipole moments of their respective bonds must be added vectorially, considering the bond angles and a bit of trigonometry. This in the beginning.
Water has one of the largest dipole moments that could be expected for a covalent compound. In the upper image we have that hydrogen atoms have positive partial charges, δ +, while oxygen carries a negative partial charge, δ-. The O-H bond is quite polar (1.5D), and there are two of them in an H moleculetwoOR.
Usually a vector is drawn that is directed from the least electronegative atom (H) to the most electronegative one (O). Although they are not drawn, on the oxygen atom there are two pairs of unshared electrons, which further “concentrate” the negative region..
Due to the angular geometry of the HtwoOr, the dipole moments add in the direction of the oxygen atom. Note that the sum of the two μO-H would give 3D (1.5 + 1.5); but it's not like that. The dipole moment of water has an experimental value of 1.85D. The effect of the near 105º angle between the H-O-H bonds is shown here..
The dipole moment of methanol is 1.69D. It is less than that of water. Therefore, the atomic masses do not have much influence on the dipole moment; but their atomic radii are. In the case of methanol, we cannot affirm that its H-O bond has a μ equal to 1.5D; since the molecular environments are different in CH3OH and HtwoOR.
That is why the length of the H-O bond in methanol would have to be measured in order to calculate μO-H. What can be said is that μO-H is larger than μC-O, since the electronegativity difference between carbon and oxygen is less than between hydrogen and oxygen..
Methanol is listed as one of the most polar solvents that can be found along with water and ammonia..
The H-N bonds are quite polar, so nitrogen, due to its higher electronegativity, attracts electrons towards itself (upper image). In addition to this, on it we have an unshared pair of electrons, which contribute their negative charges to the δ- region. Therefore, electric charges predominate on the nitrogen atom of ammonia..
Ammonia has a dipole moment of 1.42D, less than that of methanol. If both ammonia and methanol could be transformed into drafts, it would be seen that the methanol draft has more defined poles compared to the ammonia draft.
In the case of ethanol, CH3CHtwoOH, its dipole moment is very close to that of methanol, but it tends to have lower values. As there are more carbon atoms making up the δ + region, the oxygen atom representing δ- begins to lose a little of its "relative negative intensity".
Carbon dioxide has two polar bonds, C = O, with their respective dipole moments μO-C. However, as can be seen in the image above, the linear geometry of COtwo causes the two μO-C to cancel each other vectorly, even when carbon has a positive partial charge and oxygens have negative partial charges.
For this reason carbon dioxide is an apolar molecule, since μCOtwo has a value of 0D.
Both methane and carbon dioxide share something in common: they are highly symmetric molecules. In general, the more symmetric a molecule is, the smaller its dipole moment..
If we see the CH molecule4, its C-H bonds are polar, and the electrons are directed towards the carbon atom because it is slightly more electronegative. One might think that carbon would have to be a highly negative δ- region; like an eraser with its deep red center and bluish ends.
However, by dividing the CH4 in half we would get two H-C-H halves, one on the left and the other on the right, similar to the H moleculetwoO. Thus, the dipole moment resulting from adding these two μC-H would cancel out with that of the other half. And therefore, μCH4 has a value of 0D.
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