lecture b1 lewis dot structures and covalent …1. first proposed by g.n. lewis in his seminal 1916...

Lecture B1Lewis Dot Structures and

Covalent Bonding

G.N. Lewis & Linus PaulingTwo American Chemists

Linus Pauling1901-1994

G. N. Lewis1875-1946

1. First proposed by G.N. Lewis in his seminal 1916 J. Am. Chem. Soc. publication (before QM was discovered!).

2. Lewis proposed that Covalent bonds consist of shared pairs of electrons. He created a powerful empirical formalism (Lewis dot structures) for understanding bonding in simple compounds.

3. Linus Pauling created a picture of covalent bonding that employed Quantum Mechanics (and won the 1954 Nobel Prize for it).

The Covalent Bond

G.N. Lewis, circa 1902 - the cubical atom.

G.N. Lewis, Journal of the American Chemical Society 38 (1916) 762.


I2 molecule

...we still think about covalent bonds the way G.N. Lewis did. In contrast to ionic bond, electrons are shared, not transferred...

...if the atoms are identical, then electrons are equally shared - an example of a perfectly nonpolar covalent bond.

In covalent bond formation, atoms go as far as possible toward completing their octets by sharing electron pairs.Covalent bonds are drawn as lines.Lone pairs are drawn as pairs of dots.

I ⎯ I I2 molecule

Lewis dot structures provide a simple, but extremely powerful,

formalism for representing covalent bonding in molecules.

a “single bond”, consisting of a pair of electrons.



a nonbonding pair of electrons: a“lone pair”



Note: no geometric information!

Lecture B1Systematic Method for Creating

Lewis Dot Structures

Lewis Dot Structures can be produced by following a sequence of steps. Let’s produce a Lewis Dot Structure for:

NH4+ (the ammonium ion).

Step 1: Count valence electrons:

N = 54 x H = 4 x 1 = 4“+” = -1Total = 5+4-1= 8 electrons = 4 bonds and lone pairs.





Step 2: Arrange the atoms (identify a central atom, if possible).

N H

H

H

H

+






Step 3: Place electron pairs on the atoms as either bonds or lone pairs.

N H

H

H

H

+






Step 3: Place electron pairs on the atoms as either bonds or lone pairs.

Step 4: Count electrons and compare with known count. If octet rule is satisfied for all atoms and the electron count is right, you’re done.

N H

H

H

H

+

8 valence electronsaround N!

Lecture B1Lewis Dot Structures

Multiple Bonds

...another example: Let’s produce a Lewis Dot structure for: acetone: (CH3)2CO.

Step 1: Count the valence electrons:

3 x C = 12 6 x H = 6 x 1 = 6 O = 6 Total = 12+6+6= 24 electrons = 12 bonds

or lone pairs.


Step 2: Arrange the atoms: Identify a central atom. If you have C or O to choose from, go with C.

C C

H

O

C

H

H

H

H

H


Step 3: Place 12 electron pairs on the atoms.

C C

H

O

C

H

H

H

H

H


Step 3: Place 12 electron pairs on the atoms.

C C

H

O

C

H

H

H

H

H

Nine Bonds, Three Lone Pairs


Step 4: Check for the octet rule.

C C

H

O

C

H

H

H

H

H

Nine Bonds, Three Lone Pairs

oops -- only 6 e!



C C

H

O

C

H

H

H

H

H

Ten Bonds (one double bond), Two Lone Pairs

Better!

Done!



C C

H

O

C

H

H

H

H

H

This takes practice, but then becomes tedious.Getting the polyatomic arrangement can be problematic.

How about Double 07 (Nitrogen)?

Step 1: 10 valence electrons/five bonds and lone pairs



Step 2: Arrange atoms.

N N




Step 3: Fill in electron pairs and bonds.

N N




Step 3: Fill in electron pairs and bonds.

Step 4: Octet rule obeyed!

A Triple Bond!

N N

multiple bonds have higher bond dissociation energies...

A quantitative comparison of CC bonds

N2 Bond Dissociation Energy: 941 kJ mol-1

bond lengths are inversely correlated withbond energies - shorter bonds are stronger...

Double 07 is very strong!


Formal Charge

Let’s produce a Lewis Dot structure for carbon disulfide, CS2.


C = 4 2 x S = 6 x 2 = 12 Total = 4 + 12 = 16 e/8 electron pairs.




Step 2: Arrange the atoms. Go with Carbon in the center.

C SS





Step 3: Fill in the electron pairs.

C SS







C SS







C SS

Good!







C SS

BUT wait, what about??

S SC

To decide between these two candidate structures, we needto calculate the formal charges on the atoms.

SCS wins!

Less formal charge is better. Thus, SCS wins!

Carbon dioxide, CO2, is “isoelectronic” with carbon disulfide, CS2.

Nitrous oxide, N2O, is also isoelectronic with carbon disulfide, CS2.


Resonance

When multiple Lewis Dot Structures exist that are differentiated only by the positions of the electrons (the

positions of the atoms are identical), the ACTUAL structure is a weighted average of these

“resonance” structures.

We predict that the three N-O bonds in the nitrate ion are identical in length.

We predict that the two C-O bonds in the acetate ion are identical in length.

Another Example:

And Another Example: Benzene

We predict that all of the C-C bonds in benzene have the same length.

Two resonance structures for benzene...

So normally, Chemists represent benzene like this:


Exceptions to the Octet Rule

an example:

Grades are often computed using a weighted average. Suppose that homework counts 10%, quizzes 20%, and tests 70%.

If Pat has a homework grade of 92, a quiz grade of 68, and a test grade of 81, then

Pat's overall grade = (0.10)(92) + (0.20)(68) + (0.70)(81) = 79.5

In the benzene, the actual structure is the average ofthe possible resonance structures,

but this is often not the case when resonance structures with a range of formal charge values are possible. In this case, the actual structure willbe a weighted average of the available resonance structures.

A what?

How do we weight resonance structures?

An example: Which is the better Lewis Dot representation for the phosphate ion, PO43-?

We have to look at the formal charge arguments here.The structure shown at right has fewer formal charges, but...P has 10 electrons - is this ok?


If 10 electrons is okay for P, then less charge is better.



If 10 electrons is okay for P, then less charge is better.

P can use its d electrons to expand its valence shell.More than 8 electrons = "hypervalent"

What if a second double bond is added?

- formal charge on that O goes from -1 to 0- formal charge on P goes from 0 to -1 - total number of formal charges is unchanged.- electron count on P goes from 10 to 12.

If 12 electrons is okay for P, then is it better to charge P or O?

Actually, all of these resonance structures contribute to the weighted average resonance hybrid that is observed.

higher weighting factor

lower weighting factor

lower weighting factor

We need more information to determine the exact weighting factors.

Pauling will introduce the concept of electronegativity to help with this.

Final example: Rank these Lewis dot resonance structures for cyanate (NCO-) from best to worst:

Final example: Rank these Lewis dot resonance structures for cyanate (NCO-) from best to worst:

Again, Pauling's concept of electronegativity will

help with this.

1

2

3

Another exception to the octet rule:

1. Be, B, and Li can be electron deficient.

examples: note: no formal charges here...



examples: what about this resonance structure?



examples: a smaller weighting factor for this structure.+1

00

-1


2. Non metals in period 3 or higher can accommodate more than 8 electrons by employing d-orbitals.

examples:

We talked about this one before.



examples:

Sulfur can do this too.



examples:

Even Xenon gets in on the act!

Lecture B1Bond Polarity andElectronegativity

Pauling realized that covalent bonds between different nuclei will not be symmetric in electron density (probability):


This will lead to a pair of net positive and negative partialcharges that will result in a DIPOLE MOMENT for themolecule:


or

dipole moment = charge x distance


Peter J. W. Debye1884-1966

dipole moment = charge x distance

1 Debye = 1D = 3.336 x 10-30 C m

Named after Peter J. W. Debye

Dutch PhysicistNobel Prize in Chemistry 1936



HCl Dipole Moment: 1.08D

Pauling proposed a number to quantitate the ability of an atom to polarize a bond: "Pauling Electronegativity"

Electronegativity ranges from 0.7 to 4.0. It is unitless!

χ (chi)

Ionization energy: the energy required to remove an electron from an atom or ion.

Electron affinity: the energy change associated with the addition of one electron to an atom or ion.

Electronegativity: the ability of an atom in a molecule to draw bonding electrons to itself. Important: this is a dimensionless parameter - not an energy.

Three terms that are often confused:

Dipole Moments and Pauling Electronegativity



χCl - χH = 3.0 - 2.1 = 0.9

Electron density is higher on the atom with higher electronegativity.

Dipole Moments and Pauling Electronegativity


HF Dipole Moment: 1.83D

χF - χH = 4.0 - 2.1 = 1.9


χCl - χH = 3.0 - 2.1 = 0.9

HBr Dipole Moment: 0.82D

χBr - χH = 2.8 - 2.1 = 0.7

Electronegativity differencescorrelate with dipole moment values.

Rank these Lewis dot resonance structures for cyanate (NCO-) from best to worst:

We can go back and use electronegativity to rank Lewis dot resonance structures.

χN = 3.0 and χO = 3.5.

Thus the negative charge prefers to reside on the Oxygen.

1

2

3

For example, even though NH4+ has no dipole moment, the four covalent bonds in the ammonium ion are between two elements with different electronegativities: χN - χH = 3.0 - 2.1 = 0.9, and therefore EACH has a bond dipole moment.

N H

H

H

H

+

Generalizing, any bond will possess a bond dipole moment with a magnitude that is proportional to the difference in electronegativity between the bonding partners.

The fact that the four dipole moments cancel is a matter of geometric structure. That is the NEXT topic -- VSEPR Theory!

lecture b1 lewis dot structures and covalent …1. first proposed by g.n. lewis in his seminal 1916...

Documents