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Polar Covalent Bonds

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Polar Covalent Bonds• Atoms with equal or nearly equal

electronegativities form covalent bonds in whichboth atoms exert equal attractions for the bondingelectrons.

This type of covalent bond is nonpolar bond.164

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But, if the two atoms that are combined via covalentbond are different, then there is unequal sharing ofelectrons.This is due to differences in the electronegativities ofthe two atoms involved in bonding.

Electronegativity is a measure of the tendency of anatom to attract a bonding pair of electrons.

The Pauling scale is the most commonly used.Fluorine (the most electronegative element) is given avalue of 4.0, and values range down to cesium andfrancium which are the least electronegative at 0.7.

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The electronegativities of selected elements.

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• In covalent bonds with substantial difference inelectronegativity, the electrons involved inbonding are not shared equally.

• The more electronegative atom has a greaterattraction for the bonding and takes the largestshare of the electron density.

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Electronegativity of 2.1 Electronegativity of 4.0

H F

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• The result is a polar covalent bond, a bond withuneven distribution of electron density.

• The degree of polarity of a bond particularlydepends on the difference in electronegativities ofthe two atoms bonded together and partly on otherfactors such as the size of the atoms.

Such types of covalent bonds are Polar Covalentbonds.

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The distribution of electrons in a polar molecule aresymbolized by partial (d) charges.

H F

electron richregion

electron poorregion

d+ d-

Another way of representing the different electron densitieswithin a molecule is by a crossed arrow, that points from thepartially positive end of a molecule to the partially negativeend.

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Differences in electronegativities (ΔEN) between twoelements involved in chemical bonding determineswhether the bond is either ionic or covalent. If ΔEN isgreater than 2.1, the bond is ionic and if it is less than 2 itis covalent. Among covalent bonds, for non-polarmolecules ΔEN is 0, for weak polar molecules ΔEN isbetween 0.1 and 1, while strongly polar molecules ΔEN isbetween 1.1 and 2.

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The dipole moment is a measure of the polarity of amolecule. The more polar the bond the higher thedipole moment.A dipole moment is reported in a unit called a debye(D) (pronounced de-bye).

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Dipole Moment

Table: The dipole moments of some common bonds

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In a molecule with only one covalent bond, thedipole moment of the molecule is identical to thedipole moment of the bond. For example, thedipole moment of hydrogen chloride (HCl) is 1.1 D,because the dipole moment of the single bond is1.1 D.The dipole moment of a molecule with more thanone covalent bond depends on the dipole momentsof all the bonds in the molecule and the geometryof the molecule.

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chargesbetween distanceChargerQμ

• A polar bond has a dipole—it has a negative endand a positive end.

• The size of the dipole is indicated by the dipolemoment, which is given the Greek letter, m.

• The dipole moment of a bond is equal to themagnitude of the charge on the atom (either thepartial positive charge or the partial negativecharge, because they have the same magnitude)times the distance between the two charges

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The polarity (dipole moment) of a molecule is thevector sum of the dipole moment of eachindividual polar bond.For example, let’s look at the dipole moment ofcarbon dioxide. The individual carbon–oxygen bonddipole moments cancel each other—because thebond angle in CO2 is 180°—giving carbon dioxide adipole moment of zero D.

CO O

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Therefore, for molecules that have more than onecovalent bond, the geometry of the molecule must betaken into account because both the magnitude andthe direction of the individual bond dipole moments(the vector sum) determine the overall dipolemoment of the molecule. Symmetrical molecules,therefore, have no dipole moment.

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m 0 D m 1.87 D

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Bond vs. Molecular dipole moment

C CH

ClCl

HC C

Cl

HCl

H

> 0 = 0

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Unshared pairs (lone pairs) of electrons makelarge contributions to the dipole moment.

The dipole moment of water (1.85 D) is greater thanthe dipole moment of a single bond (1.5 D) becausethe dipoles of the two bonds reinforce each other.The lone-pair electrons also contribute to the dipolemoment. The dipole moment of ammonia is 1.47

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Formula (D) FORMULA (D)

H2 0 CH4 0

Cl2 0 CH3Cl 1.87

HF 1.91 CH2Cl2 1.55

HCl 1.08 CHCl3 1.02

HBr 0.80 CCl4 0

HI 0.42 NH3 1.47

BF3 0 NF3 0.24

CO2 0 H2O 1.85

Dipole Moments of Some Simple Molecules