chapter 2

Chapter 2Polar Covalent Bonds; Acids and Bases

Organic Chemistry

Chapter Objectives

Take an in-depth look at polarity of molecules

Use formal charges to designate the distribution of electrons

Represent molecules with resonance structures by ‘pushing’ electrons

Examine the acid-base behavior of molecules

Predict acid-base reactions from pKa values

Electronegativity

electronegativity – (EN) a measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound

The difference in electronegativity values for two atoms will indicate whether the two atoms form an ionic bond or a polar or nonpolar covalent bond.

Electronegativity Table

ionic bonding - 2.0 < EN polar covalent bonding - .5 ≤ EN ≤ 2.0non-polar covalent bonding - EN < .5

Bond Formation

Ionic bonding involves the loss of an electron due to a large difference in electronegativity (2.0 < EN )

Covalent bonding involves the sharing of electrons Equal sharing: non-polar bond (EN < .5) Unequal sharing: polar bond (.5 ≤ EN ≤ 2.0)

Polarity

If one side is more electronegative, it tends to have a partial negative charge (δ-) [electron-rich]

The other side tends to have a partial positive charge (δ+)[electron-poor]

The δ- and δ+ difference along a bond is called a dipole moment

δ-δ+

Ionic Character

Electrostatic Potential Maps

Red – electron rich (δ-) Blue – electron poor (δ+)


C CH

H H

HRed – electron rich (δ-) Blue – electron poor (δ+)



You describe it…What molecule do you

think it is? Take a guess…

Inductive Effect

2.1

2.1

2.1

2.12.5

3.52.5

3.5

acetic acid (ethanoic acid)(this is a carboxylic acid due to the –OH off the carbonyl group)

inductive effect – the shifting of electrons in a σ (sigma) bond in response to the electronegativity of nearby atoms.

Let’s discuss what it means to be an acid.

Inductive Effect

Why would HCN allow the H+ to be released (proton donor – acid), thus categorizing HCN as an acid, when CH4 is not usually categorized as an acid?

Dipole Moment Calculations

dipole moment (μ – Greek mu) – the magnitude of the charge (Q) at either end of the molecular dipole times the distance (r) between the charges

measured in debyes (D) μ = Q x r Just be familiar with magnitude of values

and that the D following the value is the unit.

Section 2.2

Dipole Moments

2.1

2.1

2.1

2.1

2.5

3.52.5

3.5

acetic acid (ethanoic acid)

overall dipole moment = 1.70 D

Dipole Moments

acetic acid (ethanoic acid)

overall dipole moment = 1.70 D

Water and Ammonia

Dipole Moment Values

Non-Polar Molecules

You Try It.

Draw the complete Lewis Structure for the alcohol, methanol (methyl alcohol). Show the general direction of its dipole moment.

(μ =1.70)

You Try It.

Determine if the following molecules are polar or non polar. Show any bond dipoles.

(a) (b) (c)O O

OH

You Try It.

Draw Lewis Structures for each of the following molecules and predict whether each has a dipole moment. If you expect a dipole moment, draw it in the correct direction.

(a) C2HF (b) CCl4 (c) CH3CHO

Formal Charges

formal charges – these charges do not imply the presence of actual ionic charges …instead they give insight into the distribution of electrons

calculating the formal charges of each atom in a molecule will help you determine the best, most favorable structure (lowest energy)

Section 2.3

General Rules of Stability

Lewis structures that approximate the actual molecule most closely are those that have:

maximum number of covalent bonds minimum separation of unlike charges formal charges of zero are ideal placement of any negative charges on the

most electronegative atom (or any positive charge on the most electropositive atom) Ex. Oxygen would rather 1- then 1+

Formal Charges

formal charge is calculated in the following manner:

If it violates HONC 1234, then it will have a formal charge on it.

1FC= # of valence electrons - # of non-bonding electrons+ # of bonding electrons

2

DMSO (dimethyl sulfoxide)

Formal Charges Summary

Nitromethane

Determine any formal charges on nitromethane, CH3NO2

Formal Charges

Give the formal charges for any atom on each of the following compounds Recall, having an overall + charge means that there

is one less electron

CH4 H3O+ NH3BH3

Formal Charges

Give the formal charges for any atom on each of the following compounds Recall, having an overall + charge means that there

is one less electron

H2C=N=N O3

[H2CNH2]+ (This has resonance structures.)(1 very likely, 1 less likely, 1 very

unlikely)

Resonance Structures

Some molecules cannot be represented by a single structure. In these cases we draw structures that contribute to the final

structure but which differ in the position of the bond(s) or lone pair(s).

Such a structure is delocalized and is represented by resonance forms

The resonance forms are connected by a double-headed arrow.

Section 2.4


When two or more structures are possible, the molecule will show characteristics of each structure.

Experiments show that these two structures are equivalent…both C-O bonds are same length and strength.

Since both structures are equally likely, the real structures is most likely a perfect blend of each of these. This is not always the case.


Draw resonance structures for NO3-

The “real” structure is a resonance hybrid Each oxygen has a partial negative charge

N

O

OO

_ _

N

O

OO

_

N

O

OO


The “real” structure is said to have its electrons delocalized and is represented by a dotted bond

Remember, different resonance structures are not always equivalent.

Why would the one on the left be least influential?


In some cases, one resonance form is more stable than another

(one accommodates formal charges better)

Rules for Resonance Forms

When drawing resonance structures, follow these rules:

1. Individual resonance forms are imaginary, not real2. Resonance forms differ ONLY in the placement of

their pi or non-bonding electrons3. Different resonance forms of a substance don’t

have to be equivalent4. All resonance forms must be valid Lewis

structures and obey normal rules for valency5. The resonance hybrid is more stable than any

individual resonance form

Section 2.5


1. 2. Resonance forms differ ONLY in the placement of

their pi or non-bonding electrons



1. 2.

3. Different resonance forms of a substance don’t have to be equivalent


When two resonance forms are nonequivalent, the actual structure of the resonance hybrid is closer to the more stable form than to the less stable form.



1. 2.

3.

4. All resonance forms must be valid Lewis structures and obey normal rules for valency

5.



1. 2.

3.

4.

5. The resonance hybrid is more stable than any individual resonance form


In general, the larger the number of resonance forms, the more stable (lower in energy…less reactive) a substance is because electrons are spread out over a larger part of the molecule and are closer to more nuclei.

The more the negative charge is spread out, the better!


General Trends

+ C

- C

+ N/O

- N/O

Remember

Curved arrows always represent the movement of electrons, not atoms.

Electrons always move towards the more electronegative element or positive charges.

Electron pairs can only move to adjacent positions.

The Lewis structures that result must be valid and have the same net charges.

You Try It!

Do the Resonance and Formal Charges Worksheet.

Radicals

radical - (free radical) a neutral substance that contains a single, unpaired electron in one of its orbitals, denoted by a dot (·) leaving it with an odd number of electrons.

Radicals are highly reactive! (octet rule) Radicals can form from stable molecules and

can also react with each other.

Section 2.6

Radical Resonance

Resonance forms for radicals will depend upon three-atom groupings that contain a multiple bond next to a p-orbital.

Pentadienyl Radical

You try it.

Show all the resonance forms for the straight chained C7H9

. radical (in line angle).

.

Chapter 2Polar Covalent Bonds; Acids and Bases

Part II

Organic Chemistry

Define and describe acids and bases based on the Brønsted-Lowry and Lewis definitionsUse the curved-arrow formalism to show movement of electrons between Lewis acids and basesDetermine conjugate acid-base pairsPredict strength of acids and bases based on size, electronegativity, resonance stabilization, hybridization, and induction Predict reactions using pKa values

Section 2.7-2.11 Objectives

Why Study Acids & Bases?

At a deeper level, acid-base strength allows us to predict reactivity. Compounds tend to react in such a way that

they become more stable (in the long run). Compounds considered “strong” are called that

(technically) because they dissociate completely, but (practically) also because they tend to react quickly. (This occurs because of LOW stability.)

Compounds considered “weak” tend not to react quickly or completely because they are stable. (“happy” where they are )

Everything wants to be at the

lowest possible energy.

(most stable)

Stability

Acids and Bases

Definitions of acids and bases: Brønsted-Lowry definition

Acids donate protons. (H+) (proton donor) Hint for recognizing the acid – look for Hs!

Bases accept protons. (proton acceptor) Lewis definition

Acids accept electrons. (electrophile)Bases donate electrons. (nucleophile)

Hint for recognizing the base – look for electrons!Either a lone pair or pi bonded electrons seek out

electrophiles.

Sections 2.7 & 2.11

Lewis Acids and Bases

Acid-base reactions can take place with or without a proton.

Lewis bases are species that are able to donate a pair of electrons - called nucleophiles (lover of nuclei).

Lewis acids are species that can accept this same pair of electrons – called electrophiles (lover of electrons).

The reactions are drawn using curved arrow formalism (movement of electrons represented with arrows).

Section 2.11

Group 3A elements, such as B (BF3) and Al (AlCl3), are Lewis acids because they have an unfilled valence p-orbital and can accept electron pairs from Lewis bases.

The Lewis definition of acidity includes metal cations, such as Mg2+

They accept a pair of electrons when they form a bond to a base.

Transition-metal compounds, such as TiCl4, FeCl3, ZnCl2, and SnCl4, are Lewis acids

Some Lewis Bases

Morphine – Acid and Base

Acid Reactions

So, according to Brønsted-Lowry definition, if something loses a hydrogen, it has acted as an acid. It then has the capability to accept a proton. Therefore, what is it called at this point?

A BASE! Acids will donate a proton (react) to become a

conjugate base. Bases will accept a proton (react) to become a

conjugate acid.

General Acid-Base Reaction

When writing reactions, we show the movement (called an “attack”) of electrons with an arrow. Full headed arrow – both electrons Half headed arrow – one electron

(electrophile) (nucleophile)

BE CAREFUL!

Only move electrons!!!! Never start your arrow on an atom. Start it on a pair of electrons.

…and since electrons don’t move towards other electrons, be careful where you point your arrow…Point your arrow onto an atom (NOT electrons).

Acid-Base Reactions


Strong nucleophiles are usually very strong Brønsted-Lowry bases.

Brønsted-Lowry Theory

What is the acid, base, conjugate acid, and conjugate base?

acid base conjugate conjugate acid base


Brønsted-Lowry Theory

base acid conjugate conjugate acid base

What is the acid, base, conjugate acid, and conjugate base?

(nucleophile) (electrophile)

Dual Personality

amphoteric – a substance, that depending on the circumstances, can act like an acid or a base (like water!)

What makes hydrogen sulfate ion amphoteric?

- +4 2 4

- + -24 4

HSO + H H SO

HSO H + SO

Conjugate Acid-Base Pairs

+ -2 3HCl(aq) + H O(l) H O (aq) + Cl (aq)

acid conjugate acid

base conjugate base

+ -3 2 4NH (aq) + H O(l) NH (aq) + OH (aq)

acidbase conjugate acid

conjugate base


(nucleophile) (electrophile)

Conjugate Acid-Base Pairs

The stronger the acid, the weaker the conjugate base.

The weaker the acid, the stronger the conjugate base.

The stronger the base, the weaker the conjugate acid.

The weaker the base, the stronger the conjugate acid.

You Try It

Write the acid-base reaction for each. Label the acid, base, conj. acid, conj. base. Label the electrophile and nucleophile. Show the curved arrow formalism

CH3CH2OH and NaNH2

CH3COOH and NaOCH3

CH3CH2OH and HCl

You Try It

What is the conjugate base of the following acids?

1. CH3COOH

2. CH3CH2NH3+

3.

Messing With Stability

If I take something that is stable and change it by taking something away from it, what happens? It becomes unstable. Is this good or bad?

Ex. CH3OH a weak reagent Pretty stable (How do I know this?) If I remove an “H” – CH3O-

Not stable at all a strong reagent

Acid and Base Strength

Up to this point, the terms we have been using to describe acid-base strength have been very relative.

Actual numerical values exist. Recall the Ka value

Section 2.8

The Equilibrium Expression(Law of Mass Action)

x y

eq n m

nA + mB xC + yD

[C] [D] K =

[A] [B]

The relationship between the concentration of products and reactants at equilibrium can be expressed by K

Acid Dissociation Constant

What if it is a reaction for the dissociation of an acid?

aK - +2 3

+ -3

a

HA + H O A + H O

[H O ][A ]K =

[HA]

What does the size of Ka mean?

High Ka = strong acid

Low Ka = weak acidacid dissociation constant

Acid Strength

IN CHEMICAL REACTIONS, THE ARROW USUALLY FAVORS THE PRODUCTION OF A WEAKER ACID AND BASE!!!

Why? What favors a weak acid over a strong one?

Weak acids and bases are more STABLE. If they weren’t stable, they would react to become stable…that’s why they are weak!

2 3 a

3 2 3 3 a

K

K

HCl H O H O Cl HIGH

CH OH H O H O CH O LOW

Strong acid

Weak acid Strong base

Weak base

What kind of values do you expect?

Ka vs. pKa

Acids with a greater Ka value are stronger than acids with a smaller Ka value

Problem with Ka relatively inconvenient because Ka values are usually on a negative power of ten Example: 1.0 x 10-4

To make things easier, the value pKa is used:

loga apK K

Calculating pKa

Determine the pKa of Hydrofluoric acid: Ka = 3.5 x 10-4

Phosphoric Acid: Ka = 7.5 x 10-3

~Which of these acids is stronger? H3PO4

pKa of HF:

pKa of H3PO4: What do you notice about pKa value compared to

acid strength?

3.52.1

The smaller the pKa, the stronger the acid.

pKa and Acid Strength

The smaller the pKa, the stronger the acid.

(Remember this!!!)

Table 2.3, page 51

Predicting Acid-Base Reactions from pKa Values

Do all acids react with all bases?

NO!!!!!!!!!!!!

How do we know when an acid will react with a particular base?

pKa values

Weak acids won’t produce strong acids…so, compare the pKa values of the acid and the conjugate acid to reveal which reaction proceeds, the forward or the reverse.

Section 2.9

Compare the strengths of the acid and the conjugate acid.

Predicting Acid-Base Reactions from pKa Values

You Try It

You Try It

Will the following reaction occur?

pKa = 49 pKa = 16

- -4 3 3 3CH + CH O CH + CH OH

You Try It

Write the products of the reaction and determine if it will occur.

You Try It

Predicting Acid-Base Strength without pKa values

Use the pKa values if they are handy. Otherwise…

5 major factors exist which affect acid strength1. Electronegativity2. Size3. Resonance stabilization (delocalization)4. Hybridization5. Induction

Predicting Acid-Base Strength without pKa values

1. Electronegativity

1. Electronegativity

Which will give up a hydrogen ion (proton) more readily?CH4, NH3, H2O, HF

HF is most electronegative therefore the HF bond is shared unequally and easier to break

THE MORE ELECTRONEGATIVE THE CONJUGATE BASE, THE STRONGER THE ACID

2. Size

2. Size

Which is most reactive?HF, HCl, HBr, HI Recall:

If the negative charge is spread out more, it is a more stable conjugate base…therefore…

3. Resonance Stabilization

Again…if the charge on the conjugate base is spread out more than it is a more stable conjugate base. Therefore, the original acid is a stronger acid.

…so, how does resonance help this?

The negative charge of a conjugate base may be delocalized over several atoms thus making it more stable.

3. Resonance Stabilization

The negative charge of a conjugate base may be delocalized over several atoms thus making it more stable.

You Try It Which is the strongest acid?CH3CH2OH, CH3COOH, CH3SO3H

4. Hybridization

H3C-CH3 < H2C=CH2 < HC≡CH

worst acid best acid

sp3 sp2 sp

25%-s 33%-s 50%-s The more percent “s” in character, the closer the

electrons are to the nucleus, therefore the more polarized the structure…the H becomes more positive due to the pull of e- towards the C – makes a better acid

4. Hybridization

5. Inductive Effects – e- Withdrawing

Electronegative elements “take away” electron density from a negative charge:

Stabilityincreases

5. Inductive Effects – e- Withdrawing

5. Inductive Effects – e- Donating

hyperconjugation - Donation of a pair of bonding electrons into an unfilled or partially filled orbital

5. Inductive Effects – e- Donating

Which is the most stable conjugate base?

O– O – O

–

somewhat destabilizing very destabilizing!

Molecular ModelsSection 2.12

Noncovalent Interactions

Intermolecular Forces of Attraction dipole-dipole interactions (polar molecules) London forces or dispersion forces (All molecules

have this but it is the only force present in nonpolar molecules)

hydrogen bonding (polar molecules with F, O, or N bonded to a H)

hydrophilic – water loving (attracted to water) hydrophobic – water fearing (not attracted to water)

Section 2.13

Dipole-dipole Forces

Dipole-dipole interactions are attractions between the partial negative ends and partial positive ends of two different molecules (polar molecules).

Dispersion Forces

London forces or dispersion forces are attractions due to temporary dipoles.

All molecules have this but it is the only force present in nonpolar molecules.

Hydrogen Bonding

Hydrogen bonding is the attraction between a low electronegative H on one molecule and a high electronegative F, O, or N bonded to an H on another molecule.

Hydrogen Bonding

chapter 2

Documents