Editing Chemical polarity (section)

===Polar molecules{{anchor|Polar molecules}}===<!-- This section is linked from [[Soap]] -->
[[File:Dipoli acqua.png|thumb|200px|right|The water molecule is made up of oxygen and hydrogen, with respective electronegativities of 3.44 and 2.20. The electronegativity difference polarizes each H–O bond, shifting its electrons towards the oxygen (illustrated by red arrows). These effects add as vectors to make the overall molecule polar.]]

A polar molecule has a net [[dipole]] as a result of the opposing charges (i.e. having partial positive and partial negative charges) from polar bonds arranged asymmetrically. [[Properties of water|Water]] (H<sub>2</sub>O) is an example of a polar molecule since it has a slight positive charge on one side and a slight negative charge on the other.  The dipoles do not cancel out, resulting in a net dipole. The dipole moment of water depends on its state. In the gas phase the dipole moment is ≈ 1.86 [[debye]] (D),<ref>{{cite journal |last1=Clough |first1=Shepard A. |last2=Beers |first2=Yardley |last3=Klein |first3=Gerald P. |last4=Rothman |first4=Laurence S. |title=Dipole moment of water from Stark measurements of H2O, HDO, and D2O |journal=The Journal of Chemical Physics |date=1 September 1973 |volume=59 |issue=5 |pages=2254–2259 |doi=10.1063/1.1680328|bibcode=1973JChPh..59.2254C }}</ref> whereas liquid water (≈ 2.95 D)<ref>{{cite journal |last1=Gubskaya |first1=Anna V. |last2=Kusalik |first2=Peter G. |title=The total molecular dipole moment for liquid water |journal=The Journal of Chemical Physics |date=27 August 2002 |volume=117 |issue=11 |pages=5290–5302 |doi=10.1063/1.1501122|bibcode=2002JChPh.117.5290G |doi-access=free }}</ref> and ice (≈ 3.09 D)<ref>{{cite journal |last1=Batista |first1=Enrique R. |last2=Xantheas |first2=Sotiris S. |last3=Jónsson |first3=Hannes |title=Molecular multipole moments of water molecules in ice Ih |journal=The Journal of Chemical Physics |date=15 September 1998 |volume=109 |issue=11 |pages=4546–4551 |doi=10.1063/1.477058|bibcode=1998JChPh.109.4546B }}</ref> are higher due to differing hydrogen-bonded environments. Other examples include sugars (like [[sucrose]]), which have many polar [[hydroxyl|oxygen–hydrogen]] (−OH) groups and are overall highly polar.

If the bond dipole moments of the molecule do not cancel, the molecule is polar. For example, the [[water molecule]] (H<sub>2</sub>O) contains two polar O−H bonds in a [[Bent molecular geometry|bent]] (nonlinear) geometry. The bond dipole moments do not cancel, so that the molecule forms a [[Dipoles#Molecular dipoles|molecular dipole]] with its negative pole at the oxygen and its positive pole midway between the two hydrogen atoms. In the figure each bond joins the central O atom with a negative charge (red) to an H atom with a positive charge (blue).

The [[hydrogen fluoride]], HF, molecule is polar by virtue of polar covalent bonds{{snd}}in the covalent bond electrons are displaced toward the more electronegative fluorine atom.

[[File:Ammonia-elpot-transparent-3D-balls-A.png|thumb|150px|left|The ammonia molecule, NH<sub>3</sub>, is polar as a result of its molecular geometry. The red represents partially negatively charged regions.]] [[Ammonia]], NH<sub>3</sub>, is a molecule whose three N−H bonds have only a slight polarity (toward the more electronegative nitrogen atom).  The molecule has two lone electrons in an orbital that points towards the fourth apex of an approximately regular tetrahedron, as predicted by the [[VSEPR theory]]. This orbital is not participating in covalent bonding; it is electron-rich, which results in a powerful dipole across the whole ammonia molecule.

[[File:Ozone-resonance-Lewis-2D.svg|300px|right|Resonance Lewis structures of the ozone molecule]]
In [[ozone]] (O<sub>3</sub>) molecules, the two O−O bonds are nonpolar (there is no electronegativity difference between atoms of the same element). However, the distribution of other electrons is uneven{{snd}}since the central atom has to share electrons with two other atoms, but each of the outer atoms has to share electrons with only one other atom, the central atom is more deprived of electrons than the others (the central atom has a [[formal charge]] of +1, while the outer atoms each have a formal charge of −{{1/2}}). Since the molecule has a bent geometry, the result is a dipole across the whole ozone molecule.