Chapter 2, Part I:

Polarity, Formal Charge and Resonance

Polarity and its Consequences I. Bond Polarity Polar bonds can be identified as covalent bonds in which the atoms involved have significantly different electronegativity values (EN) Organic chemists regard bonds in which Δ EN < 0.5 to be fairly nonpolar Example: C – H bonds Common polar bonds in organic compounds: somewhat polar C – N or C = N C – Cl C – Br

very polar C – O or C = O O–H C–F C – Mg or C – Li

Partial charges (δ+ and δ-) & arrow notation

II. Molecular polarity For a molecule to be polar: 1) It must contain polar bonds 2) They must be arranged unsymmetrically (shape matters) How polar is a molecule?

Dipole moment (μ) = Q (charge) x r (distance) As Δ EN and atomic radii increase, μ increases Effect of lone pairs?

III. Formal charges: An electron deficiency or excess assigned to an atom (usually when normal bonding patterns are not followed) Formal charge = number of valence e- - 1/2 the number of - number of atom normally has bonding enonbonding e-

Resonance Found in molecules in which the electron distribution can be represented in more than one way, often involving the location of bonds or nonbonding electron pairs Example: CH3NO2 Representing such molecules:

resonance forms & resonance hybrids

Although this is often used to represent benzene:

These are more accurate, since benzene is a resonance hybrid: or

An example from general chemistry: CO32O

O C O

O

C O

O

O

C O

O

Guidelines for understanding resonance forms for a molecule: 1. Resonance exists because the hybrid is more stable than any of the resonance forms. 2. The total number of electrons and the overall net charge on the molecule do not change. Only the placement of electrons and resulting formal charges on atoms may change. 3. Resonance forms must be valid Lewis structures and obey normal rules of valency. 4. Resonance forms do not have to be equivalent in stability; they may be neutral or contain atoms with charges.

Drawing resonance forms: Why? Molecules that have resonance are sometimes shown as a single resonance contributor but to understand reactivity, we must know all resonance forms They are useful for predicting regions of electron density or cationic character How? 1. Spot common features: Look for three-atom groups with double bonds and exchange bond positions

X=Y–Z*

Example: acetate ion 2. ONLY move the electrons or bonds, never move the atoms themselves! 3. Watch direction of electron movement (often depicted by curved arrows): • Move π electron pairs toward a + charge or toward another π−bonded atom. • Move nonbonding electron pairs toward an sp2-hybridized atom Example: butadiene H 2C

C H

C H

CH2

H2C

C H

C H

CH2

H2C

C H

C H

CH2

4. Obey the octet rule To avoid electron overload, don’t move more electrons toward an sp3-hybridized atom because they usually have an octet already Example: acetic acid

Chapter 2

Part II: Review of Acids & Bases Lewis Acid: Accepts an electron pair Base: Donates an electron pair

Bronsted-Lowry Acid: Donates H+ Base: Accepts H+ I.

Acid-base reactions (B-L) H–A

+

Acid II.

:B

A:-

+

Base

Conjugate base

H – B+ Conjugate acid

Acid strength depends on acid dissociation constant, Ka HA Ka

+ =

pKa =

H2O

A-

+

H3O+

[H3O+] [A-]

Keq [H2O]

=

- log Ka

Trend: The greater the pKa, the weaker the acid

[HA]

III.

Acid/base behavior and structure of organic compounds Functional groups Inductive effect of electronegative atoms Stability of conjugate base

IV.

Predicting the direction of equilibrium Recall that a strong acid forms a weak conjugate base and vice versa In acid-base equilibrium, reaction favors formation of weaker acid or base

V.

Lewis acid-base reactions: electrons from the base are shared to form the new bond Electron flow can be depicted by using curved arrows: Example: H3C

H NH2

H

B

H H

H3C

N H

H B H

H

Acid and Base review (cont’d) 1) Try to think of acids and bases more from a Lewis perspective: Lewis bases are electron donors: they contain electron-rich atoms such as O or N with lone pairs that can donate the lone pair for bonding Lewis acids are electron acceptors: they may be electron deficient with less than an octet or positively charged like H+ or carbocations. 2) Draw arrows from the bonding e- of the Lewis base to the target. Those electrons should form a covalent bond. 3) Lewis acid-base reactions may form only one product or more than one depending on the nature of the donor and acceptor 4) Do not randomly break other bonds that are not part of the base-to-acid transfer of electrons. Cl H3C

N H

CH3

Cl

Al Cl

OH H

Br

H NO 3

H3C

H3C

C H

O

..

CH3

CH3

F H3C

N H

CH3

F

B F

Some characteristic organic acids and bases Carboxylic acids behave like acids (thus the name) and donate H+ Ex:

Acetic acid, pKa = 4.76 CH3COOH +

H2O

Weak acid

CH3COO- +

H3O+

Equilibrium favors reactants

strong conj. base

Amines generally behave like bases and accept protons Ex:

Methylamine (pKa = 40) CH3NH2

+

Weak base

CH3COOH weak acid

CH3NH3+

+

conj. acid

CH3COOconj. base

Alcohols are considered neutral, but during a reaction can accept or donate a H+ Ex:

Methanol (pKa = 15.5)

In the presence of base, methanol loses a proton: CH3OH

+

Weak acid

NaNH2

Na CH3O-

+

base

strong base

NH3 strong conj. acid

In the presence of acid, methanol can be protonated CH3OH

+

Weak base

CH3COOH

CH3OH2+

weak acid

conj. acid

+

CH3COOconj. base

Following the flow of electrons: Reaction mechanisms with curved arrows Some rules to follow in the use of curved arrows 1.

Always draw arrows so they point in the direction electrons are moving: from a pair of electrons in the base to the acidic H:

2.

You may need to move more than one e- pair in a given step…remember, no atom can be left with more than an octet! If an atom with a full octet accepts an e- pair, it must donate another one.

3.

Indicate any changes in charge that take place (be careful of +/- signs) a) With negatively charged bases, the atom donating the e- pair for bond formation becomes neutral after the new bond forms, so the conjugate acid is neutral b) If the base is neutral, after donating the e- pair it becomes a positively charged conjugate acid

4.

Never use arrows to show where atoms are going, only where electrons are going!

5.

In most reactions, electrons move in pairs, so use a double-headed arrow For free radical reactions, movement of a single electron is shown by a “fishhook”

6.

The curved-arrow rules also apply to other types of reactions, not just traditional acid-base reactions. Electron or bond donors are called “nucleophiles” Electron acceptors are called “electrophiles”

Electronegativity chart of the elements Yellow = electropositive Green = intermediate

Polarity of bonds affects dipole moment of molecules

Red = electronegative

To calculate formal charges correctly, first draw the Lewis structure

Acid-base chemistry

pKa = 40

Representing structure of organic compounds

Skeletal or line-bond structures do not show the H atoms

Resonance forms and electron flow