Zumdahl’s Chapter 8

Chemical Bonding:

Stealing and Sharing

of electrons

Chapter Contents

• Types of Bonds

• Electronegativity

• Polarity & Dipoles

• Ions

• Binary Ionic Compounds

• Polar Covalency

• The Covalent Bond

• Bond Energies

• Local & Global Model

• G.N. Lewis Structures

• Octet or not Octet

• Resonance fudge

• Valence Shell Electron Pair Repulsion Theory

Types of Chemical Bonds

• Ionic Bonds• Large differences in electronegativity hold

atoms together by Coulombic potentials.

• Radically polar poorly-shared binding electrons

• Covalent Bonds• Near equally-shared, weakly polar binding pair

• Dative Bonds (donor-receptor model)

• Metal Bonding (macroscopic wavefunction)

Electronegativity,

• Measure of ionization potential and electron affinity. The power to hold and attract.• Empirically, deviation of bond energies, DHX,

from the geometrical average of DHH and DXX.

= |A – B| determines ionicity of AB < ~1 “covalent” while > ~2 “ionic”

• “Polar covalent” for between 1 and 2.• Not yet ionic, but still significantly polar.

Pauling’s Definition

• DAB = [ DAADBB ]½ + 96.5( )²

H2.1

PaulingElectronegativities

He

Li1.0

Be1.5

B2.0

C2.5

N3.0

O3.5

F4.0

Ne

Na0.9

Mg1.2

Al1.5

Si1.8

P2.1

S2.5

Cl3.0

Arsoon

K0.8

Ca1.0

Ga1.6

Ge1.8

As2.0

Se2.4

Br2.8

Kr3.0

Rb0.8

Sr1.0

In1.7

Sn1.8

Sb1.9

Te2.1

I2.5

Xe2.7

Cs0.7

Ba0.9

Tl1.8

Pb1.9

Bi1.9

Po2.0

At2.2

Rn

Dipole Moment, , & Polarity

0 implies charge separation in bonds.

• Charge separation means bond polarity:• Product of Q times RAB, bond distance.

• Molecular dipole is vector sum of bond dipoles.• Symmetries cause vector cancellations!

• Bond polarity is necessary but not sufficient cause of molecular polarity.

Ions

• Electron configuration follows removal of electrons with highest n from parent atom.

• Effective charge governs size of ion:• Cations are much smaller than parent atom.

• Anions are much larger but• High charge anions are unstable except in

compounds because 2nd extra electron is repelled by 1st anion’s negative charge.

Recipe for a Crystal

• Gasify both the electropositive and the electronegative elements. ( Hphase )

• Dissociate both to atoms. ( DXX )

• Make the cation and anion. ( IP & EA )

• Condense the crystal from the infinitely remote ions. ( Ulattice )

• Add + and – components Hformation

Madlung Constants

• Lattice energies can be calculated by crystal geometries. Nearest neighbors attract, next nearest neighbors repel, and so on forever.

• Madlung constants (k = 1.5– 4.2) sum those all up as a scale factor for binary ionic potentials.

• While O2– can’t exist by itself (O– has a positive electron

affinity), it is stabilized in Mg2+O2–, say, by its very negative lattice energy, 4 that of NaCl due to Q1Q2, with which it shares a crystal structure (same k).

Polar Covalency

• Most compounds have heteronuclear bonds, so 0 and the bond is polar.• But few will have so high (>~2) that their

bonds are truly ionic.

• Slight inequality in electron sharing makes them “polar covalent.”

• % ionic = 100% [ measured / ionic ]

• Not to be confused with ionic % in .

Covalent Bonds

• Electron pairs shared (more or less equally) between bonding partners.

• Build-up of e– density between nuclei.• Counteracts nuclear (proton) repulsion

• Increases attractions (to neighboring protons)

• Stretches ; helps minimize e– kinetic energy

• Reduces e– density near nucleus, reducing the e––e– repulsions there.

Two H atoms NOT Bonding

Notice theextent of overlap.

H2 with Bonding turned ON

DeeperOverlap

“Bond Energies”

• While it takes a definite energy to break a particular bond in a unique molecule, “bond energies” refers to some weighted average of a particular bond in all the molecules in which it is found.• As such, they are correct for no molecule.

• Still energies can be estimated from – D, for bonds changed during a reaction.

Bond Enthalpy

H ~ – D since H measures differences in the formation of compounds from elements while D shows the differences in the destruction of compounds. Opposites.• So H ~ D broken – D formed

• If D formed > D broken , products are more stable than reactants, and we expect exothermic reaction, heat evolved as potential lowers.

Covalency is Local

• In the Local Electron model, electrons are• Shared in pairs between adjacent atoms,

• unshared entirely as lone pairs,

• or hidden deep and uninvolved as core.

• This ignores the reality of global electrons, free to roam over more than one atom pair.• G.N. Lewis had a patch for this difficulty with

the LE model, but we need a better model.

Gilbert Newton Lewis

• Developed a method to manipulate valence electrons to satisfy local atomic “needs.”• Devised Lewis Structures and rules for the

placement of electron pairs about them to• Put at least one pair between all bonded atoms.

• Complete “rare gas electronic configurations” about all the molecule’s atoms.

• Arrange multiple bonds (multi-electron pairs) between atoms to minimize “formal charges.”

Lewis Structure Rules

• Sum all valence and ionic electrons = N.

• Pick an atomic skeletal structure.

• Place two of N between all bonded atoms.

• Distribute remainder as lone pairs to achieve an octet (only duet for hydrogen).

• Minimize formal charge by stealing lone pairs to make additional (multiple) bond pairs. If FC is dumb, pick another skeleton.

Formal Charges

• Shared electrons count toward BOTH atoms’ octets.

• But shared electrons are divided equally between their bonded atoms for FC.

• Lone pairs count fully in FC of an atom.

• FC = # valence electrons – sum of above.

• Best value is FC=0 for all atoms.• But FC = ion’s charge, so some won’t be 0.

• Negative FC goes to highest electronegativity.

Egregious Example, NOCl

• Nitrosyl chloride, “NOCl,” is a horrific non-aqueous solvent for some food processing applications; O=N–Cl seems more likely.

• Try to find a good NOCl Lewis Structure:• # = 5+6+7 = 18 valence electrons

• N–O –Cl uses 4, so use 7 lone pairs (14 e–)

• :N:–:O:–:Cl:: fails to deliver a N octet

• :N:=O:–:Cl:: gives everyone an octet; OK?

NOCl Bombs in Formal Charge!

• Since N, O, and Cl are expecting 5, 6, and 7 valence electrons, respectively,• :N:=O:–:Cl:: then shows “formal charges” of

• –1 +1 0• putting the positive charge on the most

electronegative atom?!?

• It’s more likely we had the skeleton wrong, so let’s try ONCl.

ONCl, a Better Lewisite

• O–N–Cl still needs 7 lone pairs

• ::O:–N:–:Cl:: still needs N octet help

• ::O=N:–:Cl:: satisfies octets & gives FC of

• 0 0 0 Perfect!

• So Lewis Structures permit us to correct an incorrect molecular formula into one truly reflecting the geometry, ClNO.

Octet Trumps Formal Charge

• If you can’t get both, sacrifice good formal charges to securing an octet. For example,• CO has 4+6=10 valence electrons

• C–O needs 4 more lone pairs

• :C:–:O: gives nobody an octet

• :C O: isoelectronic with :NN:, but FC are

• –1 +1 not what one would expect at all!• but that inobvious polarity is correct.

Breaking the Octet

• “Rare Gas Configuration” is 8 valence electrons (save for He) ending in ns2 np6.

• Atoms beyond row 2 have d orbitals which empower them to adopt more than 8.

• Often “central,” such atoms can surround themselves with typically 12 but sometimes 14 or more valence electrons in their molecules.

Meta Arsenic Acid

• (HO)AsO2 has As as its central atom.

• # = 5+3(6)+1 = 24 valence electrons

• ::O = As = O:: distributes all 24

• | and puts octets on all oxygens

• ::O – H (and the duet on hydrogen)

• with zero formal charges everywhere,

• but requires 10 not 8 electrons on As.

• HNO3 exists too but has a N octet and FCs!

The Impoverished 2nd Row

• But perchloric acid, (HO)ClO3

• ::O – H

• |

• ::O = Cl = O::

• ||

• :O:

• has no fluorine (octet) analogue.• No 2d orbitals & F can’t take +3 formal charge!

Newest Rare Gas Compound

• HArF (sounds like a hairball), recently synthesized, leaves only He and Ne as truly noble gases.

• But neither has d orbitals.

• It’s Ar’s d orbitals that permit H–Ar:::–F::: with its 10 electrons around Ar.

• Don’t hold your breath for Ne compounds.

Resonance

• Often several different legitimate Lewis Structures are possible. NO3

– ion is all of

• : O:: _ : O : _ : O:: _

• | and || and |

• ::O=N–O::: :::O–N–O::: :::O–N=O::

• with FCs of +1 (N), –1 (O–), and 0 (O=)

• representing delocalized electrons, greater stability, and uniform NO bond orders of 1.33!

Valence Shell Electron Pair Repulsion Theory (VSEPR)

• Lewis Structures show e– pairs on & between atoms.

• VSEPR uses those to predict shapes as a consequence of e––e– repulsions such that• Lone pairs repel more strongly than bond pairs.

• All pairs seek geometries that distance them.

• Pairs in multiple bonds are treated as if they are a single bond pair for directional purposes.

• Nuclei follow blindly where bonding e–s point.

# of directional electron pairs:

• 1: linear A-B molecule

• 2: collinear• 180°

• 3: trigonal• 120°

• 4: tetrahedral• 109.47°

• 5: trigonal bipyramid• 120° and 90°

• 6: octahedral• 90°

Molecular Shapes

• Determined only by the direction of the bonding pairs since only they terminate in atoms.

• Lone pairs still dictate where bonding pairs go, but lone pair directions aren’t involved in describing molecular shapes.

• A Xn Em used to count bonding directions (n) and lone pairs (m). A is central atom.

Examples of 5 Directional Pairs

• AX5

• Trigonal bipyramid

• PCl5

• AX4E

• See Saw

• SF4

• AX3E2

• T-shaped

• BrF3

• AX2E3

• Linear

• I3–

Central Atoms

• … aren’t terminal atoms in the molecule.

• Molecules may have several central atoms.

• Geometry is determined at each by VSEPR.

Aceticacid

Histamine Nitroglycerin

Mint