ElectronegativityElectronegativity

+ – 0 0

H Cl H H

The basic units: ionic vs. covalentThe basic units: ionic vs. covalent• Ionic compounds form repeating units.• Covalent compounds form distinct molecules.• Consider adding to NaCl(s) vs. H2O(s):

HO

H Cl Na

Na Cl

Cl

Cl

Na

Na H

O H

HO H

• NaCl: atoms of Cl and Na can add individually forming a compound with million of atoms.

• H2O: O and H cannot add individually, instead molecules of H2O form the basic unit.

I’m not stealing, I’m sharing unequallyI’m not stealing, I’m sharing unequally• We described ionic bonds as stealing electrons• In fact, all bonds share – equally or unequally.• Note how bonding electrons spend their time:

• Point: the bonding electrons are shared in each compound, but are not always shared equally.

• The greek symbol indicates “partial charge”.

H2 HCl LiCl

+ –0 0 + –

covalent (non-polar) polar covalent ionic

H H H Cl [Li]+[ Cl ]–

ElectronegativityElectronegativity

• Recall that electronegativity is “a number that describes the relative ability of an atom, when bonded, to attract electrons”.

• The periodic table has electronegativity values.

• We can determine the nature of a bond based on EN (electronegativity difference).

EN = higher EN – lower ENNBr3: EN = 3.0 – 2.8 = 0.2 (for all 3 bonds).

Bond Type

Basically: a EN below 0.5 = non-polar covalent, 0.5 - 1.7 = polar covalent and above 1.7 = ionic

Determine the EN and bond type for these: HCl, CrO, Br2, H2O, CH4, KCl

Electronegativity AnswersElectronegativity Answers

HCl: 3.0 – 2.1 = 0.9 polar covalent

CrO: 3.5 – 1.6 = 1.9 ionic

Br2: 2.8 – 2.8 = 0 covalent

H2O: 3.5 – 2.1 = 1.4 polar covalent

CH4: 2.5 – 2.1 = 0.4 covalent

KCl: 3.0 – 0.8 = 2.2 ionic

Holding it togetherHolding it togetherQ: Consider a glass of water. Why do molecules

of water stay together?A: there must be attractive forces.

Intramolecular forces occur between atoms

Intermolecular forces occur between molecules

• We do not consider intermolecular forces in ionic bonding because there are no molecules.

• The type of intramolecular bond determines the type of intermolecular force.

Intramolecular forces are much

stronger

Inter-molecular Forces (IMF’s)

• London Dispersion Forces

• Dipole – Dipole Forces

• Hydrogen Bonding

London Dispersion Forces

• London dispersion forces result from instantaneous non-permanent dipoles created by random electron motion. London dispersion forces are present in all molecules and are directly proportional to molecular size.

• LDF’s are the weakest of the IMF’s

Dipole – Dipole Forces

• Dipole-dipole interaction is the attraction between a partially negative portion of one molecule and a partially positive portion of a nearby molecule.

• Dipole-dipole interaction occurs in any polar molecule as determined by molecular geometry.

Hydrogen Bonding

• Hydrogen bonding is the unusually strong dipole-dipole interaction that occurs when a highly electronegative atom (N, O, or F) is bonded to a hydrogen atom.

• Hydrogen bonding is the strongest of the 3 IMF’s

Effects of IMF’s on Physical Properties

• The strength of intermolecular forces present in a substance is related to the boiling point and melting point of the substance. Stronger intermolecular forces cause higher melting and boiling points.

EXAMPLES:

• CH4 - Methane: has only very weak London dispersion forces (lowest b.p. & m.p.)

• CHCl3 - Chloroform: has dipole-dipole interaction (moderate b.p. & m.p.)

• NH3 - Ammonia: has hydrogen bonding and dipole-dipole interaction (high b.p. & m.p.)

1. PbS2. HClO4

3. Zn(OH)2

4. HBr(aq)5. (NH4)2CO3

6. sulfur hexafluoride7. nitrous acid8. hydrogen chloride9. lead(II) chloride10. zinc sulfate

Name / Formulate the following Compounds

1. Lead (II) Sulfide2. Hydrogen Perchlorate3. Zinc Hydroxide4. Hydro Bromic Acid5. Ammonium Carbonate6. SF67. HNO2(aq)8. HCl9. PbCl210. ZnSO4