CHEM 200/202Professor Jing GuOffice: EIS-210

All emails are to be sent to:chem200@mail.sdsu.edu

My office hours will be held in online on Monday 9-11 AM or by appointment.

ANNOUNCEMENTS• Class (https://SDSU.zoom.us/j/281477051)

• Go to the help room GMCS-212 (https://SDSU.zoom.us/s/239381878)

• Go to the Supplemental Instruction sessions(http://bit.ly/virtualCHEM200)

• The exam 2 will be postponed

PROBLEM• Are the bonds in each of the following

substances ionic, non-polar covalent, or polar covalent? Arrange the substances with polar covalent bonds in order of increasing bond polarity.

• S8

• RbCl

• PF3

• SCl2

• F2

• SF2

• Rank the members of each set of compounds in order of increasing ionic character of their bonds. Use polar arrows to indicate the bond polarity of each.

• HBr, HCl, HI

• H2O, CH4, HF

• SCl2, PCl3. SiCl4

QUESTIONArrange the bonds below in order of increasing ionic

character, and include the direction of the bond polarity with a polarity arrow.

C－BC－O C－C C－NC－F

Most IonicLeast Ionic

QUESTIONArrange the bonds below in order of increasing ionic

character, and include the direction of the bond polarity with a polarity arrow.

C－B C－OC－C C－N C－F

Most IonicLeast Ionic

∆EN= 0.00 0.49 0.51 0.89 1.45

LEWIS ELECTRON-DOT SYMBOLS

• The Lewis dot system is a means of representing the valence electrons and the bonds formed by atoms.

• For main group elements, the group number identifies the number of valence electrons.

• Each dot represents a valence electron; dots are placed around the “nucleus” on all four sides, and only become paired when more than four electrons are present.

Nitrogen5 valence electrons N•••••

N• •• •

•N•• •••

N••• ••

QUESTIONWhich is the correct Lewis electron-dot

symbol for carbon?

C••• •

•C• ••C• ••• •

C• •••

Answer:

- A- B- C

- D

LEWIS ELECTRON-DOT SYMBOLS

Dots represent the valence electrons around the symbol of the element. No more than one pair of electrons per

side, pairing only occurs on the 5th electron.

CHEMICAL VALENCE:

THE NUMBER OF VALENCE ELECTRONS OF AN ATOM WILL INFLUENCE HOW MANY BONDS IT MAY FORM.

C• ••• N• •••

• O• •••• • F••

••• ••

Group 14chem. valence 4

Group 15chem. valence 3

Group 16chem. valence 2

Group 17chem. valence 1

CH H

H

H

NH H

H

••OH H••••

FH••••

••

FOLLOWING THE OCTET RULE

A total of four electron pairs, either bonding or lone pairs, will be associated with each atom.

X X•• X

X X•• X

••••

X X••

X••••

X••••

••

CONVERTING FORMULAS INTO LEWIS STRUCTURES

1. Start with molecular formula.

2. Place atoms - typically atom with least electronegative atom in center.

3. Valence electrons - sum the total # of valence e- for all atoms

4. Draw single bonds between atoms - subtract 2e- for each single bond, 4e- for double bonds...

5. Ensure each atom has 8e- (2e- for H) - use lone pairs.

LEWIS STRUCTURES

• Hydrogen atoms are terminal.

• Halogen atoms are terminal. Exceptions: the heavier halides (Cl, Br, I) are central atoms only if they bind to more than one O or F atom.

• The less electronegative atoms will be the central atom.

• More than one central atom: the less electronegative atom will be surrounded by more atoms than the more electronegative atom.

• Structures with the maximum number of bonds are preferred.

The following rules allow us to predict the most stable Lewis structure:

LEWIS STRUCTURES• Draw the Lewis structures for the following compounds:

• Methane (CH4)

• Water (H2O)

• Carbon dioxide (CO2)

• Ethanol (CH3CH2OH)

• Phosgene (PH3)

HYPERVALENT LEWIS STRUCTURES

• Atoms from period 3 and beyond (n≥3) can have an expanded octet (more than 8 electrons) due to their potential to have d orbitals.

• Draw the Lewis structures for the following compounds:

• Sulfuric acid (H2SO4)

• Phosphoric acid (H3PO4)

• Perchloric acid (HClO4)

• Phosphorus pentachloride (PCl5)

• Boron trifluoride (BF3)

LECTURE OBJECTIVES• Chapter 7.4

• Compute formal charges on atoms in a Lewis structure.

• Use formal charges to identify the most stable isomer or resonance Lewis structures.

• Explain the concept of, and draw, Lewis structures.

ISOMERSIsomers are compounds with identical molecular formulas but different bonds forming the compound.

Example: Molecular formula CHN

H C N••

Hydrogen cyanideH N C ••

Hydrogen isocyanide

Relative Stability of IsomersMore stable isomers have:

• the less electronegative atoms as central atoms• atoms at their expected valency• minimized formal charges on atoms

ASSIGNMENT OF VALENCE ELECTRONS IN COMPOUNDS

• Formal Charge

• The formal charge is the charge on an atom in a compound assuming perfect covalent bonding.

• Shared electrons (covalent bond electrons) are divided equally between bonded atoms.

FORMAL CHARGE & ISOMERS

H C N•• H N C ••

Hydrogen cyanide Hydrogen isocyanide

FormalCharge 0 0 0 0 +1 -1

The most stable isomer is the form as the formal charges on all atoms is zero.

Carbon: 4(ve) - 0(lp) - 4(bp) = 0Nitrogen: 5(ve) - 2(lp) - 3(bp) = 0

Carbon: 4(ve) - 2(lp) - 3(bp) = -1Nitrogen: 5(ve) - 0(lp) - 4(bp) = +1

Formal Charge = # of valence e-s - # lone pair e-s - 1e- per bond pair

FORMAL CHARGE CALCULATION

C

O

O O

•• ••

•• ••••

••••••

2-

Formal Charge = # of valence e-s - # lone pair e-s - 1e- per bond pair

Formal chargesCarbon = 4(ve) - 0(lp) - 4(bp) = 0Oxygen = 6(ve) - 4(lp) - 2(bp) = 0Oxygen = 6(ve) - 6(lp) - 1(bp) = -1Oxygen = 6(ve) - 6(lp) - 1(bp) = -1

OXYGEN-OXYGEN BONDSBy examining known oxygen bonds we can see that the

predicted Lewis structure of ozone doesn’t match the reality.

H O•• O H•••• ••

O•• O•••• ••

Hydrogenperoxide

Bond length (pm) Bond energy (kJ)

149 204

Oxygen 121 498

Ozone O•• O•••• O•••••• 128 310

(same for both bonds)If the Lewis structure was correct each of the two bonds would have different BE and lengths, similar to those of

oxygen and hydrogen peroxide.

RESONANCEElectrons are not static points on structures; they are dynamic and

constantly in motion.

The third O-O bond is said to be delocalized - the electrons are distributed between all three oxygen atoms.

O•• O•••• O••••xxx

x

O•• O•••• O•••••• O••O•• ••O••••

••Resonance

structures for Ozone:

Delocalized bonding electrons

Delocalized lone pairResonance Hybrid:

RESONANCE & BOND ORDER

O•• O•••• O••••xxx

x

O•• O•••• O•••••• O••O•• ••O••••

••Resonance

structures for Ozone:

Resonance structures can have non-integer bond orders.

O-O Bond order = (3 e- pairs)÷(2 resonance structures)O-O Bond order = 1½

JUDGING THE RELATIVE IMPORTANCE OF RESONANCE STRUCTURES

•O3 and other compounds (e.g. NO3-, CO32-...) have identical atoms surrounding the central atom - so all resonance structures have identical formal charges.

•Compounds where the atoms around the central atom are not identical give rise to different resonance structures which contribute to the resonance hybrid to differing extents.

O N N•••••• •• O•• N•• N•••• O•• N•• N•• ••

Less important More importantResonance

Hybrid for N2O: O•• N•• N••xx

xx

FormalCharge +1 +1 -2 0 +1 -1 -1 +1 0

QUESTION

Which is the most important resonance structure for the thiocyanate ion [NCO]-?

Structure AStructure BStructure CAll three are equalBoth A & C are better than B

Answer:ABCDE

O C N•••••• •• O•• C•• N•••• O•• C•• N•• ••[ ]- [ ]- [ ]-

A B C

RESONANCE

FormalCharge +1 0 -2 0 0 -1 -1 0 0

The B and C structures have the formal charges closest to 0 on most atoms.

Structure C has the negative formal charge on oxygen, which is favored as it is the most electronegative atom of the compound.

Structure C contributes most to the resonance hybrid structure.

O C N•••••• •• O•• C•• N•••• O•• C•• N•• ••[ ]- [ ]- [ ]-

A B C

RESONANCE• Selection criteria for the most important resonance structure:

• Smaller formal charges (positive or negative) are preferable.

• Avoid charges of the same polarity on adjacent atoms (-- or ++).

• If a negative formal charge is required it should be on the atom with the greatest electronegativity value.

ELECTRON DEFICIENT AND ODD ELECTRON LEWIS STRUCTURES

• Draw the Lewis structures for the following compounds:

• Boron trihydride (BH3)

• Nitric oxide (NO)

• Nitrogen dioxide (NO2)