CHEM 200/202Prof. Christopher R. Harrison

Office: GMCS-213E

Office hours via Zoom at https://sdsu.zoom.us/j/

99415148959Mon. & Wed. 12:00-1:00

Dozer

ANNOUNCEMENTS• Exam results have posted - review them.

• Go to the help room to figure out what you got wrong on the exam.

• Science Communication Essay #2 due Oct. 22nd.

• Chapter 7-9.3 homework due Nov. 12th. - do the intro to ChemDoodle before starting!

LECTURE OBJECTIVES• Chapter 7.3

• Write Lewis symbols for neutral atoms and ions.

• Draw Lewis structures depicting the bonding in simple molecules.

• 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.

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.

1)HBr, HCl, HI

2)H2O, CH4, HF

3)SCl2, PCl3. SiCl4

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 VALENCEThe number of valence electrons an atom has will influence how many bonds it may form.

Group Valence Electrons

Chemical Valence

Lewis Structure

Possible Bonds

14 4 4 4

15 5 3 3

16 6 2 2

17 7 1 1

C• •••

N• ••••

O• •••• •

F•••

•• ••

CHEMICAL VALENCE

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••••

••

LEWIS THEORY OF CHEMICAL BONDING• Lewis structures - represent valence electrons around the

atom; only pair electrons when more than 4 valance electrons.

• A covalent bond in a molecule forms when two atoms in the molecule share one, or more, electron pairs.

• Octet Rule: when main-group elements form bonds, they lose, gain or share valence electrons so as to “have” a total of eight.

• Lewis structures depict the connectivity of the molecule.

• Classifies valence electrons in the molecule:

• Bonding electrons - shared Lone pairs - unshared

CO O••••

••••

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:

• Oxygen (O2)

• Water (H2O)

• Carbon dioxide (CO2)

• Phosgene (PH3)

• Ethanol (CH3CH2OH)

• Propene (C3H6)

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 (BrF3)

ELECTRON DEFICIENT AND ODD ELECTRON LEWIS STRUCTURES

• Draw the Lewis structures for the following compounds:

• Boron trihydride (BH3)

• Nitric oxide (NO)

• Nitrogen dioxide (NO2)

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

Example: Molecular formula CHN

H C N••Hydrogen cyanide

H 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

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

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 charges

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.