WS 3:Dipole Moment1 In a hydrogen molecule, each nucleus has a unit charge of +1. Electrons are attracted equally to both nuclei. The result is a symmetrical orbital in which electron density near one nucleus is the same as the electron density near the other nucleus. Although a more complicated atom, when fluorine bonds to itself, a similar situation occurs for the formation of the fluorine molecule. In the chemical bond between the two identical atoms, the bonding electrons are shared equally. In contrast to a symmetrical molecule, an unsymmetrical molecule has unsymmetrical attractive forces. The bonding electrons are attracted by different nuclear charges. On the hydrogen atom, we see a +1 nucleus, while on F, there is a +9 nucleus, screened by electrons in 1s, 2s, and 2p orbitals. We have already discussed the shielding abilities of s and p orbitals, so we know the bonding pair don’t feel a full +9 charge, but they feel a positive charge greater than the charge on the hydrogen atom. Unsymmetrical attractive forces lead to an unsymmetrical distribution of electrons. The electron density is concentrated closest to the nucleus with the larger Zeff. The bond is polarized because of this shift in density. The hydrogen is a little positive because it has lost some electron density and the fluorine is more negative because it has gained electron density. This is a polar-covalent bond. Polar covalent bonds are said to arise when there is an electronegativity difference between atoms. This difference is sufficient to cause partial charges on the bonding atoms, but not so large that it produces an ionic bond.

The dipole moment The difference in electronegativity leads to the polar covalent bond. The dipole moment is a measurable quantity of the polarity of the bond. The symbol for the dipole moment is µ. It is a vector quantity that has both magnitude and direction. The formula for dipole moment is: µ = q X r, where r is the distance between the two charges and q is the magnitude of the charges. When we look at the dipole moment we often look at systems that have equal charges that are opposite in sign. The dipole moment is measured in Debyes (1D=3.34X10- 30C•m). The dipole is often represented as an arrow with a plus on its tail. The plus sign represents the more electropositive nucleus while the arrow tip points to the more electronegative atom. [pages 313-315 in your textbook].

this can also be written as:

δ+ δ− H — Cl 1 BackgroundH— readingCl for this handout can be found in section 8.4 of Brown and Lemay. ws 3- dipole moment.docx Page 1 of 3 Alscher

Questions To Start You Thinking: Which has the larger dipole moment in each of the following cases? In group 1 r ≠ R. in group 2, r=r. A B +1 –1 +1 –1

r R

A B +0.5 –0.5 +1 –1

r r

1. Consider HF and HCl:

a) Which is the bigger atom, F or Cl?

b) Which has the longer bond length r, HF, or HCl?

c) Which has the greater partial charge F in HF or Cl in HCl? (hint, which is more electronegative?)

d) Why is the dipole moment of HF larger than the dipole moment of HCl? Another way of wording this is: which factor seems more important for determining the dipole moment: the bond length or the partial charge?

ws 3- dipole moment.docx Page 2 of 3 Alscher 2. Base on the dipole moments for HF [1.82 D], HCl [1.08D], HBr [0.83D], HI [0.45D], which is more important in determining the dipole moment—bond length or the electronegativity difference?

3. The iodine monobromide molecule, IBr, has a bond length of 2.49 Å and a diopole moment of 1.21 D. (a) Which atom of the molecule is expected to have a negative charge? Explain. (b) calculate the effective charges on the I and Br atoms in IBr in units for the electronic charge, e.

4. Use vector notation to designate the dipole moments of CO, HI, and ClF.

ws 3- dipole moment.docx Page 3 of 3 Alscher