chapter 2 polar covalent bonds; acids and bases part i organic chemistry

Chapter 2Polar Covalent Bonds; Acids and Bases

Part I

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 – 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.

Bond Formation

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

Covalent bonding involves the sharing of electrons Equal sharing: non-polar bond (EN<.5) Unequal sharing: polar bond (.4<EN<2.1)

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+

Electrostatic Potential Maps

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

Electrostatic Potential Maps

C CH

H H

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

Electrostatic Potential Maps

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

Electrostatic Potential Maps

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

You describe it…What molecule do you

think it is? Take a guess…

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

Dipole Moment Calculations

Section 2.2 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

Dipole Moment Values

Inductive Effect

2.1

2.1

2.1

2.12.5

3.52.5

3.5

acetic acid (ethanoic acid)

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

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?

You Try It.

Draw the complete Lewis Structure for the alcohol, methanol (methyl alcohol). Show the direction of its dipole moment. (μ =1.70)

You Try It.

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

(a) (b) (c)O O

OH

You Try It.

Draw a Lewis Structure of 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 (Section 2.3)

formal charges – these charges don’t 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)

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+

DMSO (dimethyl sulfoxide)

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

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

CH3NO2 H2C=N=N O3

[H2CNH2]+ (draw all resonance structures)(1 very likely, 1 less likely, 1 very

unlikely)

Resonance Structures

formaldehyde

Resonance Structures

Some compounds are not adequately represented by a single Lewis structure as we saw in the previous example.

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

Resonance Structures

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

Resonance Structures

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

Resonance Structures

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

(one accommodates formal charges better)

Resonance Structures

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

Resonance Structures

When drawing resonance structures, follow these rules:

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

their pi or non-bonding electrons

Resonance Structures

When drawing resonance structures, follow these rules:

1. 2.

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

Resonance Structures

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

General Trends

+ C

- C

+ N/O

- N/O

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.

Radical Resonance

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

Pentadienyl Radical

Pentadienyl Radical

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 & 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

Morphine

Acid Reactions

So, 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 to become a

conjugate base. Bases will accept a proton 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

Acid-Base Reactions

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?

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

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 between

CH3CH2OH and NaNH2

Write the acid-base reaction between

CH3COOH and NaOCH3

Write the acid-base reaction between 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’s 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-Base Strength

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

Actual numerical values exist Recall the Ka value

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.

Reactivity

Do all acids react with all bases?

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

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

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

Use the pKa values if they are handy. Otherwise… 5 major factors exist which affect acid strength

Electronegativity Size Resonance stabilization (delocalization) Hybridization Induction

Predicting Acid/Base Strength

Electronegativity

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

Size

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…

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.

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

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

Hybridization

Inductive Effects – e- Withdrawing

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

Stabilityincreases

Inductive Effects – e- Withdrawing

Inductive Effects – e- Donating

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

Inductive Effects – e- Donating

Which is the most stable conjugate base?

O– O – O

–

somewhat destabilizing very destabilizing!

Lewis Acids & 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)

Drawn using curved arrow formalism (movement of electrons represented with arrows)

Lewis Acids & Bases

Strong nucleophiles are usually very strong Brønsted-Lowry bases (HIGH pKa) (Unstable)

Section 2.13 Noncovalent Interactions

Intermolecular Forces of Attraction dipole-dipole interactions (polar molecules) hydrogen bonding (polar molecules with F, O, or N

bonded to a H) London forces or dispersion forces (All molecules

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

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