Molecules with zero dipole moment are non-polar in nature, while molecules with dipole moment are said to be polar in nature. As the magnitude of dipole moment increases, more will be the polar nature of the bond in a molecule. The dipole moment is used to find the polar nature of the bond. The resultant dipole moment of \(NH_3\) is 1.49 D. On the other hand, \(NH_3\) has a pyramidal structure, with 3 N – H bonds and a lone pair on the nitrogen atom. The 3 bonds are present in a single plane, so dipole moments cancel each other and the net dipole moment equal to zero. Boron trihydride \(BH_3\) has a symmetrical structure and the 3 B – H bonds are at an angle of \(120^o\) to each other. The dipole moment of ammonia \(NH_3\) is 1.49D. In the boron trihydride \(BH_3\) molecule, the dipole moment is zero. The net dipole moment in a water \(H_2O\) molecule is 1.84 D. The individual bond moment of an oxygen-hydrogen bond in a water molecule is 1.5 D. The bond angle in a water \(H_2O\) molecule is \(104.5^o\). Thus, the individual bond dipole moments do not cancel each other out as in the case in the \(BeF_2\) molecule. The presence of a lone pair of electrons in the oxygen atom causes the water molecule to have a bent shape as per the Valence shell electron pair repulsion theory. In a water molecule, the electrons are localized around the oxygen atom as it is much more electronegative than the hydrogen atom. As they are equal in magnitude but are opposite in direction the net dipole moment of a \(BeF_2\) molecule is zero. The two individual bond dipole moments cancel each other in a \(BeF_2\) molecule. Fluorine is a more electronegative atom than Beryllium so it shifts the electron density towards itself. The bond angle between the two beryllium-fluorine bonds in the beryllium fluoride molecule is \(180^o\). This movement of electrons can be shown via the bond dipole moment. When two atoms with varying electronegativities interact, the electrons tend to move from their initial positions to come closer to the more electronegative atom. In chemistry, the arrows that are drawn in order to symbolize dipole moments begin at the positive charge and end at the negative charge. The direction is parallel to the bond axis. The bond dipole moment \(\mu\) is also a vector quantity. \(\delta\) is the magnitude of the partial charges \(\delta^ and \delta^- \) and d is the distance between \(\delta^ and \delta^-\). The bond dipole moment arises between two atoms of different electronegativities. The dipole moment is measured in Debye units. ![]() \(Dipole Moment (\mu) = Charge (Q) \times distance of separation (r)\) In the case of a polyatomic molecule, the dipole moment of the molecule is the vector sum of all bond dipoles in the molecule.Ī dipole moment is the product of the magnitude of the charge and the distance between the centres of the positive and negative charges in a system. ![]() This arrow shows the shift of electron density in the molecule. The dipole moment is denoted by a cross on the positive centre and an arrowhead on the negative centre.
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