2.Electronegativity and Polarity

(Silberberg: Chapters 8 & 9)

INORGANIC CHEMISTRY ICHM474

Electronegativity, Bond Polarity & Dipole Moment

Outline:

• Trends of Electronegativity

• How to determine bond polarity

• How to identify dipole moment

Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond.

Electron Affinity - measurable, Cl is highest

Electronegativity - relative, F is highest

X (g) + e- X-(g)

Electronegativity is defined as the ability of an atom in a molecule to attract electrons to itself

Prof. Linus PaulingNobel Prize for Chemistry 1954Nobel Prize for Peace 1962

Electronegativity is a function of two properties of isolated atoms;The atom’s ionization energy (how strongly an atom holds onto its own electrons)The atom’s electron affinity (how strongly the atom attracts other electrons)

For example, an element which has:A large (negative) electron affinityA high ionization (always endothermic, or positive for neutral atoms)

Will: Attract electrons from other atoms and Resist having electrons attracted away

Such atoms will be highly electronegative

Electronegativity decreases down a group for representative elements.

Electronegativity generally increases left to right across a period.

• When two identical atoms form a covalent bond each atom has an equal share of the bond’s electron pair

• When different kinds of atoms combine, one of the nuclei usually attracts the electrons in the bond more strongly

(a) The electron density of the electron pair in the bond is spread evenly between the two H atoms in H2, which gives a nonpolar covalent bond.

(b) The electron density of the bond in HCl is pulled more tightly around the Cl end of the molecule giving a partial separation of charge and a polar covalent bond.

hydrogen chlorine iodine nitrogen

Covalent bonding with equal sharing of electrons occurs in diatomic molecules formed from one element.

A dash may replace a pair of dots.

• If the two atoms that constitute a covalent bond are identical, then there is equal sharing of electrons.

• This is called nonpolar covalent bonding.

• Ionic bonding and nonpolar covalent bonding represent two extremes.

• If the two atoms that constitute a covalent bond are not identical, then there is unequal sharing of electrons.

• This is called polar covalent bonding.

• One atom assumes a partial positive charge and the other atom assumes a partial negative charge.– This charge difference is a result of the

unequal attractions the atoms have for their shared electron pair.

:H Cl+ -

Shared electron pair.

:The shared electron pair is closer to chlorine than to hydrogen.

Partial positive charge on hydrogen.

Partial negative charge on chlorine.

Chlorine has a greater attraction for the shared electron pair than hydrogen.

Polar Covalent Bonding in HCl

The attractive force that an atom of an element has for shared electrons in a molecule or a polyatomic ion is known as its electronegativity.

The polarity of a bond is determined by the difference in electronegativity values of the atoms forming the bond.

Linus Pauling devised a method for calculating electronegativities of elements.

Pauling's electronegativity values for representative elements are given in the following figure.

However, the transition metals do not follow these trends.

Nonmetals are the most electronegative elements and metals are the least electronegative (they are electropositive)

H2.1

Li1.0

Be1.5

B2.0

C2.5

N3.0

O3.5

F4.0

Na0.9

Mg1.2

Al1.5

Si1.8

P2.1

S2.5

Cl3.0

K0.8

Ca1.0

Pauling's electronegativity values of the first twenty elements.

The difference in the electronegativity values of two bonded atoms gives an estimation of the polarity to be expected in a bond.

Electronegativity values are useful in determining if a bond is to be classified as nonpolar covalent, polar covalent or ionic. What you should do is look only at the two atoms in a given bond. Calculate the difference between their electronegativity values. Only the absolute difference is important.

The three major types of intramolecular bond can be described by the electronegativity difference:

Non-Polar Covalent – Bonds which occur between atoms with

little or no electronegativity difference (less than 0.5).

Polar Covalent – Bonds which occur between atoms with a

definite electronegativity difference (between 0.5 and 2.0).

Ionic – Bonds which occur between atoms with a large

electronegativity difference (2.0 or greater), where electron transfer

can occur.

E.g. F-F (4.0 – 4.0 = 0) is non-polar covalentH-F (4.0 – 2.1 = 1.9) is polar covalentLiF (4.0 – 1.0 = 3.0) is ionic

H F+ -

I. Nonpolar Covalent:

This type of bond occurs when there is equal sharing (between the two atoms) of the electrons in the bond. Molecules such as Cl2, H2

and F2 are the usual examples.

Textbooks typically use a maximum difference of 0.2 - 0.5 to indicate nonpolar covalent. Since textbooks vary, let us use 0.5.

One interesting example molecule is CS2. This molecule has

nonpolar bonds. Since the electronegativities of C and S are both 2.5, you have a nonpolar bond.

II. Polar Covalent:

This type of bond occurs when there is unequal sharing (between the two atoms) of the electrons in the bond. Molecules such as NH3 and

H2O are the usual examples.

The typical rule is that bonds with an electronegativity difference less than 1.6 are considered polar. (Some textbooks or web sites use 1.7.) Obviously there is a wide range in bond polarity, with the differences in the C-H bonds in CH4 being 0.4 to the difference the H-O bonds in

water being 1.4.

III. Ionic:

This type of bond occurs when there is complete transfer (between the two atoms) of the electrons in the bond. Substances such as NaCl and MgCl2 are the usual examples.

The rule is that when the electronegativity difference is greater than 2.0, the bond is considered ionic.

So, let's review the rules: 1.If the electronegativity difference (usually called EN) is less than 0.5, then the bond is nonpolar covalent.

2.If the EN is between 0.5 and 1.6, the bond is considered polar covalent

3.If the EN is greater than 2.0, the the bond is ionic.

That, of course, leaves us with a problem. What about the gap between 1.6 and 2.0? So, rule #4 is:

4.If the EN is between 1.6 and 2.0 and if a metal is involved, then the bond is considered ionic. If only nonmetals are involved, the bond is considered polar covalent. So, that means compounds like HF and SiO2 are considered to be polar

covalent, even though there is a large electronegativity difference.

A warning: rule #4 may not exist in your textbook. Often, the 1.6 value is used and the 1.6-2.0 range is lumped into the ionic category. (Steven Zumdahl in his "World of Chemistry" textbook, makes the rule be 2.0 instead of 1.6 for polar covalent. This allows him to include HF as polar covalent.)

Covalent

share e-

Polar Covalent

partial transfer of e-

Ionic

transfer e-

Increasing difference in electronegativity

Classification of bonds by difference in electronegativity

Difference Bond Type

0 Covalent

2 Ionic

0 < and <2 Polar Covalent

EN

3.0

2.0

0.0

Boundary ranges for classifying ionic character of chemical bonds.

H H

Hydrogen Molecule

If the electronegativities are the same, the bond is nonpolar covalent and the electrons are shared equally.

The molecule is nonpolar covalent.

Electronegativity2.1


11.10

Electronegativity Difference = 0.0

If the electronegativities are the same, the bond is nonpolar covalent and the electrons are shared equally.

Cl Cl

Chlorine Molecule



The molecule is nonpolar covalent.


11.10

If the electronegativities are not the same, the bond is polar covalent and the electrons are shared unequally.

H Cl

Hydrogen Chloride Molecule



The molecule is polar covalent.

+ -


11.10

Sodium Chloride

Na+ Cl-

If the electronegativities are very different, the bond is ionic and the electrons are transferred to the more electronegative atom.



The bond is ionic.No molecule exists.


11.10

A dipole is a molecule that is electrically asymmetrical, causing it to be oppositely charged at two points.

A dipole can be written as + -

An arrow can be used to indicate a dipole.

The arrow points to the negative end of the dipole.

H Cl H Br H

O

H

Molecules of HCl, HBr and H2O are polar .

A molecule containing different kinds of atoms may or may not be polar depending on its shape.

The carbon dioxide molecule is nonpolar because its carbon-oxygen dipoles cancel each other by acting in opposite directions.

• The magnitude of the polarity is expressed in terms of the dipole moment

• Dipole moments are frequently reported in units of Debye (D)

chargesbetween distance

charge ofamount

moment dipole

r

q

rq

m C 103.34D 1 -30

Prof. Peter DebyeNoble Prize 1936

• The dipole moments and bond lengths for some diatomic molecules are:

115 0.16 NO

113 0.11 CO

161 0.45 HI

141 0.82 HBr

127 1.09 HCl

91.7 1.83 HF

Length(pm) Moment(D) Compound

Bond Dipole

11.11

Relating Bond Type to Electronegativity Difference.

• The difference in electronegativity provides an estimate for the degree of polarity (or sometimes referred to as the ionic character) of the bond

• There is no sharp dividing line between ionic and covalent bonding: ionic bonding and nonpolar covalent bonding represent the extremes

• A bond is mostly ionic when the electronegativity difference between the two atoms is large

• The degree of polarity, or ionic character, varies continuously with the electronegativity difference

• In general, electronegativity increases bottom to top in a group and left to right in a period

Each atom in a bond has a partial charge of about +0.5 or –0.5 units when the electronegativity difference is 1.7.

OH

H

dipole momentpolar molecule

S

OO

CO O

no dipole momentnonpolar molecule

dipole momentpolar molecule

C

H

H

HH

no dipole momentnonpolar molecule

Which of the following molecules have a dipole moment?H2O, CO2, SO2, and CH4

2.Electronegativity and Polarity

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