Slide 1

8.4 Bond Polarity and Electronegativity.

The concept of electronegativity was developed by Linus Pauling. Electronegativity is the ability of an element to attract electrons to itself in a molecule. Electronegativity increases across the periodic table and is at a maximum in the top right hand corner at fluorine, and is at a minimum at the bottom left hand corner at Cesium.

Linus Carl Pauling (1901-1994)

ELECTRONEGATIVITY

Pauling originally developed the concept from the fact that for ionic compounds the bond energies were much larger for e.g. HF, than expected from the average of the energies for the related homonuclear diatomics, in this case H2 and F2. The more the observed energy of bond formation exceeded the average of the energies of the two related homo-nuclear diatomics, the greater the electronegativity.

Covalent electron Polar covalent Ionic one atomDensity spread equally one atom has morehas attracted most ofOver both atoms electron densitythe electron density

e-densityhigh

e-densitylow

e-densityspreadequally

Electronegativities Li to F

On the next slide we have a table of electronegativities for elements in the periodic table. One sees that F (EN = 4.0) is the most electronegative element while Cs is the least electronegative (EN = 0.7) The electronegativites increase across the periodic table from Li (EN = 1.0) to Li by 0.5 per element, so that we have:

LiBeBCNOF EN: 1.01.52.02.53.03.54.0

Electronegativities of the Elements

F with EN = 4.0 ismost electronegativeelement

Cs (EN = 0.7) is least electronegativeelement

Au is at the peak ofan island of electronegativity,and is most electronegative metal

Electronegativities of some main group elements:

H2.1LiBeBCNOF 1.01.52.02.53.03.54.0NaMgAlSiPSCl0.91.21.51.82.12.53.0KCaGaGeAsSeBr0.81.01.61.82.02.42.8RbSrInSnSbTeI0.81.01.71.81.92.12.5

Electronegativity and bonding:

Electronegativity tells us what kind of bonding we have, i.e. whether it is ionic or covalent. The greater the difference in EN between the two elements forming the bond, the more ionic is the bond. Typical ranges for EN differences are:EN differencebonding typeExample EN difference range__________________________________________________________________________________> 2.0IonicLiF4.0-1.0 = 3.00.5-2.0polar covalentHF4.0-2.1 = 1.9 2.0) :

FeCoNiCuZnGaGeEN:1.81.91.91.91.61.61.8RuRhPdAgCdInSnEN:2.22.22.22.11.71.71.8OsIrPtAuHgTlPbEN:2.22.22.22.52.12.01.9

The metals with EN > 2.0 have special chemistry where they can form stable covalent bonds to carbon, for example, and have chemistry that is much more covalent than found for other less electronegative metals.

The remarkable chemistry of the metallic elements with EN > 2.0

Elements such as Pt, Ag, Au, and Hg are extremely covalent in their bonding. Thus, they form stable complexes with bonds to carbon atoms, and other elements with EN values of about 2-2.5. Examples are [Au(CN)2]- and [Hg(CN)2] (CN- = cyanide) or [Au(CH3)2]- and [Hg(CH3)2].

Hg

Structure of[Hg(CH3)2]

The inert pair:

The elements after gold in the periodic table have as their most stable oxidation state one which is 2 less than the group valency. Thus, Pb has as its most stable oxidation state the Pb(II) state, although Pb is in group 4. This is referred to as the inert pair, and is thought to be due to increased electronegativity caused by relativisitic effects. The inert pair of electrons is usually stereochemically active, as are the lone pairs on molecules such as ammonia, as expected from VSEPR:

Pb

lone pair

Cl

Structure of[PbCl3]-

The lead-acid battery works on the greater
stability of Pb(II) than Pb(IV) plus Pb(0)

The reaction at the anode involves oxidation of Pb toPbSO4(s) and at the cathode reduction of PbO2 to PbSO4(s).

anode (Pb metal)positive

cathode (PbO2)(negative)

electrolyte =dilute H2SO4

cathode (PbO2)

anode (Pb metal)

Pb(IV) + Pb(0) 2 Pb(II)

ventcasing

cell divider

cell connectors

vent caps