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Dipole moments
Polar diatomic molecules The symmetrical electron distribution in the bond of a homonuclear diatomic renders the bond non-polar. In a heteronuclear diatomic, the electron withdrawing powers of the two atoms may be different, and the bonding electrons are drawn closer towards the more electronegative atom. The bond is polar and possesses an electric dipole moment (µ). Be careful to distinguish between electric and magnetic dipole moments.
The dipole moment of a diatomic XY is given by equation 1.1 where d is the distance between the point electronic charges (i.e. the internuclear separation), e is the charge on the electron (1.602 × 10-19 C) and q is point charge. The SI unit of µ is the coulomb metre (C m) but for convenience, tends to be given in units of debyes (D) where 1D = 3.336 × 10 -30Cm.
µ = q × e × d (1.1)
Worked example: Dipole moments
The dipole moment of a gas phase HBr molecule is 0.827 D. Determine the charge distribution in this diatomic if the bond distance is 141.5 pm. (1D = 3.336 × 10-30Cm) To find the charge distribution we need to find q using the expression:
In worked example above the result indicates that the electron distribution in HBr is such that effectively 0.123 electrons have been transferred from H to Br. The partial charge separation in a polar diatomic molecule can be represented by use of the symbols δ+ and δ- assigned to the appropriate nuclear centres, and an arrow represents the direction in which the dipole moment acts. By SI convention, the arrow points from the δ- end of the bond to the end, which is contrary to long-established chemical practice. This is shown for HF in structure 1.14. Keep in mind that a dipole moment is a vector quantity.