For other uses of this term, see dipole moment.
It is defined as chemical dipole moment (μ) to measure the intensity of the attraction force between two atoms. It is the expression of the asymmetry of the electrical charge in a chemical bond. It is defined as the product between the distance "d" that separates the charges (link length) and the value of equal and opposite charges in a chemical bond:
μ = q ⋅ d {\displaystyle \mu \,=q\cdot d}
Usually it is expressed in Debyes (1 D = 1 A. 1 ues). The value of q can be interpreted as the degree of sharing of the load, that is, according to the differences in electronegativity, what percentage (100q) of the charge shared by the covalent bond is displaced towards the load in question. In other words, it represents what part of an electron is being "felt" more or less by the charges in question. Vector μ Μ partial vectors are canceled creating an apolar molecule.
The participation of the dipole moment as a vector with direction towards the most electronegative atom of the link is very important. Take the example of carbon dioxide (CO2), due to the difference of electronegativities in the C-O bonds we find a μ different from 0, but the CO2 molecule, experimentally, proves not to be polar. This is because the molecular geometry of the CO2 determines that both μ vectors of the two C-O bonds are canceled by vector addition. This shows that the polarity of the molecules depends both on the dipole moments of bonding, and on the molecular geometry determined by the TREPEV, which results in a molecular dipole moment (vector sum of the partial dipole moments). Importance of μ in the dipole-dipole attraction Main article: Dipole-dipole interaction
The intensity of the attraction forces between polar molecules (dipole-dipole) is determined by the dipole moment of each of them, acting in a directly proportional way. Thus, the more polar the molecule (higher μ), the greater the intensity of the intermolecular attraction forces of the dipole-dipole type present in the substances.
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