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Chapter 9: Problem 37

What is the distinction between a bond dipole and a molecular dipole moment?

Short Answer

Expert verified
A bond dipole is the electric dipole moment within a single bond, resulting from the difference in electronegativity between two bonded atoms, while a molecular dipole moment is the overall dipole moment of the entire molecule, obtained as the vector sum of all individual bond dipoles. Bond dipoles exist due to charge separation within a bond, while molecular dipole moments depend on both the bond dipoles and the molecular geometry, which may result in a net zero molecular dipole moment even when individual bond dipoles are present.

Step by step solution

01

Definition of bond dipole

A bond dipole is created when two atoms in a molecule have a difference in electronegativity. The more electronegative atom attracts the electrons in the bond, creating a partial negative charge, denoted as "δ-", while the less electronegative atom has a partial positive charge, represented by "δ+". This charge separation leads to the formation of an electric dipole, with the dipole moment \(\vec{\mu}\) pointing from the positive to the negative charge. It is represented as: \[ \vec{\mu} = \delta q \cdot \vec{d} \] Where \(\delta q\) is the magnitude of the partial charge and \(\vec{d}\) is the vector distance between the atoms.
02

Definition of molecular dipole moment

The molecular dipole moment is the vector sum of all bond dipoles in the molecule. In other words, it is the net dipole moment of the entire molecule, considering the geometry of the molecule and the distribution of partial charges. It is calculated as: \[ \vec{\mu}_{molecule} = \sum_{i} \vec{\mu}_{i} \] Where \(\vec{\mu}_{i}\) represents each bond dipole in the molecule.
03

Distinction between bond dipole and molecular dipole moment

Bond dipole refers to the electric dipole moment that exists within a single bond, resulting from the difference in electronegativity between the two bonded atoms. In contrast, the molecular dipole moment is the overall dipole moment of the entire molecule, which is the vector sum of all individual bond dipoles. A molecular dipole moment may end up being null if the bond dipoles present in the molecule cancel each other out due to the molecular geometry, even if individual bond dipoles exist.

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Most popular questions from this chapter

(a) If you combine two atomic orbitals on two different atoms to make a new orbital, is this a hybrid orbital or a molecular orbital? (b) If you combine two atomic orbitals on one atom to make a new orbital, is this a hybrid orbital or a molecular orbital? (c) Does the Pauli exclusion principle (Section 6.7\()\) apply to MOs? Explain.

(a) What does the term paramagnetism mean? (b) How can one determine experimentally whether a substance is paramagnetic? (c) Which of the following ions would you expect to be paramagnetic: \(\mathrm{O}_{2}^{+}, \mathrm{N}_{2}^{2-}, \mathrm{Li}_{2}^{+}, \mathrm{O}_{2}^{2-} ?\) For those ions that are paramagnetic, determine the number of unpaired electrons.

Write the electron configuration for the first excited state for \(\mathrm{N}_{2}\) - that is, the state with the highest-energy electron moved to the next available energy level. (a) Is the nitrogen in its first excited state diamagnetic or paramagnetic? (b) Is the \(\mathrm{N}-\mathrm{N}\) bond strength in the first excited state stronger or weaker compared to that in the ground state? Explain.

Consider the \(\mathrm{H}_{2}^{+}\) ion. (a) Sketch the molecular orbitals of the ion and draw its energy-level diagram. (b) How many electrons are there in the \(\mathrm{H}_{2}^{+}\) ion? (c) Draw the electron configuration of the ion in terms of its MOs. (d) What is the bond order in \(\mathrm{H}_{2}{ }^{+}\) ? (e) Suppose that the ion is excited by light so that an electron moves from a lower-energy to a higherenergy MO. Would you expect the excited-state \(\mathrm{H}_{2}^{+}\) ion to be stable or to fall apart? Explain.

The \(\mathrm{O}-\mathrm{H}\) bond lengths in the water molecule \(\left(\mathrm{H}_{2} \mathrm{O}\right)\) are \(0.96 \AA\), and the \(\mathrm{H}-\mathrm{O}-\mathrm{H}\) angle is \(104.5^{\circ} .\) The dipole moment of the water molecule is \(1.85 \mathrm{D} .\) (a) In what directions do the bond dipoles of the \(\mathrm{O}-\mathrm{H}\) bonds point? In what direction does the dipole moment vector of the water molecule point? (b) Calculate the magnitude of the bond dipole of the \(\mathrm{O}-\mathrm{H}\) bonds. (Note: You will need to use vector addition to do this.) (c) Compare your answer from part (b) to the dipole moments of the hydrogen halides (Table 8.3). Is your answer in accord with the relative electronegativity of oxygen?
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