Carbon–oxygen and hydrogen–oxygen bonds are polar because of a
difference in electronegativity of the two elements connected by the bond.
Recall that bond polarity is a measure of how equally the electrons in a bond
are shared between the two atoms of the bond. As the difference in
electronegativity between the two atoms of a bond increases, so does the
bond polarity. This polarity can be represented by a vector (called bond
dipoles).
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However, the overall polarity of a molecule is the sum of its bond
polarity (dipoles). (a) In CO2 the bond dipoles are equal in
magnitude but exactly oppose each other. The overall dipole
moment is zero. (b) In H2O the bond dipoles are also equal in
magnitude but do not exactly oppose each other. The molecule has
a nonzero overall dipole moment or the molecule is polar.
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Figure 9.10 shows examples of polar and nonpolar molecules, all
of which have polar bonds. The molecules in which the central
atom is symmetrically surrounded by identical atoms (BF3 and
CCl4) are nonpolar. For ABn molecules in which all the B atoms
are the same, certain symmetrical geometries—linear (AB2),
trigonal planar (AB3), tetrahedral and square planar (AB4), trigonal
bipyramidal (AB5), and octahedral (AB6)—must lead to nonpolar
molecules even though the individual bonds might be polar.