CHAPTER 3ACIDS AND BASES IN ORGANIC CHEMISTRY

3.1 INRODUCTION

The important of acid-base reaction:1. is a simple, fundamental reaction.2. enable you to see the mechanism of the reaction3. illustrate the process of bond breaking and bond making 4. examine important ideas about the relationship between the structure of molecules and their reactivity.5. illustrate the important role solvents play in chemical reactions6. find something familiar in General chemistry.

3.1A THE BRNSTED-LOWRY DEFINATION OF ACIDS AND BASES

According to the brnsted-Lowry theory, an acid is a substance that can donate a proton, and a base is a substance that can accept a proton. For example:

H O

H

H Cl H O+

H

H Cl-++

Base(protonacceptor)

Acid(protonsonor)

Conjugate acid of H2O

Conjugate base of HCl

Hydrogen chloride, a very strong acid, transfers its proton to water. Water acts as a base and accepts the proton.

The molecule or ion forms when an acid loses its proton is calledthe conjugate base of that acid. Such as the chloride ion. The molecule or ion that forms when a base accepts a proton is called the conjugate acid of that base. Such as the hydronium ion.

Other strong acids that completely transfer a proton when dissolvedin water are hydrogen iodide, hydrogen bromide, and sulfuric acid.

HI + H2O

HBr + H2O

H2SO4 + H2O

HSO4- + H2O

H3O+ + I-

H3O+ +Br-

H3O+ +HSO4-

H3O+ + SO42-

The proton transfer is stepwise in sulfuric acid, the first proton Transfer completely, the second only to the extent of ～ 10%.

When an aqueous solution of sodium hydroxide is mixed with an aqueous solution of hydrogen chloride, the reaction that occurs isbetween hydronium and hydroxide ions.

H O-H O+

H

H Cl-+ + Na+ + H O

H

+ Na+ + Cl-2

Spectator ions

Total Ionic Reaction:

Net Reaction:

H O+

H

H + H O- H O

H

2

The net reaction is simply:

H3O+ + OH- 2 H2O

3.1B THE LEWS DEFINITION OF ACIDS AND BASES

Lewis proposed that acids be defined as electron-pair acceptors andbases be defined as electro-pair donors. For example:

H+ + NH3 NH3+H

Lewis acid(electron-pairacceptor)

Lewis base(electron-pairdonor)

+ NH3

Lewis acid(electron-pairacceptor)

Lewis base(electron-pairdonor)

Cl

Al

Cl

Cl

Cl

Al-

Cl

Cl NH3+

+ NH3

Lewis acid(electron-pairacceptor)

Lewis base(electron-pairdonor)

F

B

F

F

F

B-

F

F NH3+

The Lewis theory , by virtue of its broader definition of acids, allows acid-base theory to include all of the Brnsted-Lowry reactions aqnd , as we shall see, a great many others.

Any electron-deficient atom can act as a Lewis acid. Many compounds containing group IIIA elements such as aluminium are Lewis acids because group IIIA atoms have only a sextet of electrons in their outer shell. Many other compounds that have atoms with orbitals also act as Lewis acids. For example:

R O H R O+

HLewis acid(electron-pairacceptor)

Lewis base(electron-pairdonor)

ZnCl2 ZnCl2-+

Br Br + FeBr3 Br Br+ FeBr3-

Lewis base(electron-pairdonor)

Lewis acid(electron-pairacceptor)

3.2 THE USE OF CURVED ARROWS IN ILLUSTRATING REACTIONS

The curved arrows is commonly used by organic chemists to showthe direction of electron flow in a reaction. Besides it is a useful Method for indicating which bonds form and which bonds break.

Now, lets us illustrate some of the basic ideas of the curved-arrownotation with simple Lewis acid-base reactions:

H O

H

H Cl H O+

H

H Cl-+ +

this bond breaks this bond is formed

The following acid-base reactions gives other examples of the use of the curved-arrow notation:

H O+

H

H O+ H- H O

H

2

Acid Base

H O+

H

H+ H O

H

Acid Base

H3C C

O

O H H3C C

O

O +

H O H

Acid Base

H3C C

O

O H H3C C

O

O +-

O+ H-

3.3 THE STRENGTH OF ACIDS AND BASES: Ka AND pKa

When acetic acid dissolves in water, the following reaction dosenot proceed to completion:

H3C C

O

O H H3C C

O

O+ H2O + H3O+

Experiments show that in a 0.1M solution of acetic acid at 25℃ onlyabout 1% of the acetic acid molecules ionize by transferring theirprotons to water. The reaction is a equilibrium, we can describe it with an expression for the equilibrium constant.

For dilute aqueous solutions, the concentration of water is essentiallyconstant, so the equilibrium constant can be expressed with the acidity constant (Ka).

Keq [H2O] =[H3O+] [CH3CO2

-] [CH3CO2H]Ka =

Keq =[H3O+] [CH3CO2

-][CH3CO2H] [H2O]

3.3A THE ACIDITY CONATANT Ka

3.3B ACIDITY AND pKa

A large value of Ka means the acid is a strong acid, and a small

value of Ka means the acid is a weaker acid. If the Ka is greater

than 10, the acid will be completely dissociated in water.

Chemists usually express the acidity constant, Ka, as its negative logarithm, pKa.

pKa = - logKa

For acetic acid the pKa is 4.75:

Notice that there is an inverse relationship between the magnitude of the pKa and the strength of the acid:

pKa = -log(1.76 ¡Á 10-5) = - (-4.76) = 4.76

The larger the value of the pKa, the weaker is the acid.

CH3CO2H ©‚ CF3CO2H ©‚ HCl

pKa = 4.76 pKa = 0 pKa = -7

Weak acid Very strong acid

increasing acid strength

In pure water at 25¡æ, the concentrations of hyfronium and hydroxideIons are equal to 10-7M. Since the concentration of water in pure water is 55.5 M, we can calculate the Ka for water

Ka =[H3O+] [OH-] [H2O] =

(10-7)(10-7)

(55.5)= 1.8 ¡Á 10-16

pKa = 15.7

3.3C PREDICTING THE STRENGTH OF BASES

The principle which allows us to estimate the strengths of bases:the stronger the acid, the weaker will be its conjugated base.

So relate the strength of a base to the pKa of its conjugate acid. Thelarger the pKa of the conjugate acid, the stronger is the base.

increasing base strength

Cl- CH3CO2- OH-

Very weak basepKa of conjugateacid (HCl) = - 7

pKa of conjugateacid (CH3CO2H) = 4.75

Strong basepKa of conjugateacid (H2O) = 15.7

Amines are like ammonia in that they are weak bases. Dissolving inwater bring about the following equilibrium.

H O H O H-NH3 + NH4+ +

Base Acid Conjugate acidpKa = 9.2

Conjugate base

Dissolving methylamine in water causes the establishment of asimilar equilibrium.

+ H O H O H-+

Base Acid Conjugate acidpKa = 10.6

Conjugate base

CH3NH2 CH3NH3+

3.4 PREDICTING THE OUTCOME OF ACID-BASE REACTIONS

There is a general order of a acidity and basity for some of thecommon acids and base. For example: acetic acid has a pKa = 4.76and carboxylic acids generally have pKa values near 3 ～ 5. The pKaof ethyl alcohol is 16, and alcohols generally have pKa values near15 ～ 18, and so on. But there are exceptions.

The general principle of predicting whether or not an acid-base will occur: acid-base reaction always favor the formation of the weakeracid and the weaker base.

Using the principle, we can predict that a carboxylic acid will react with aqueous NaOH in the following way because the reaction willlead to the formation of the weaker acid and the weaker base.

+R C

O

O H R C

O

O + H O HO H-Na+ NaH+

Stronger acid pKa = 3-5

Stronger base Weaker base Weaker acid pKa = 15.7

Because of the acidity water-insoluble carboxylic acids dissolve in aqueous sodium hydroxide; they do so by reacting to form water-soluble salts.

+C

O

O H C

O

O + H O HO H-Na+ NaH+

insoluble in water soluble in water

The water-insoluble amines dissolve readily in hydrochloric acidbecause the acid-base reaction convert them to soluble salts.

3.5 THE RELATIONSHIP BETWEEN STRUCTURE AND ACIDIY

NH2 NH3+Cl- + H O H

water insoluble soluble-soluble salt

H O+

H

HCl-+

The strength of the bond to the proton is the dominating effect if wecompare compounds in a vertical column of the periodic table.

H—F H—Cl H—Br H—IpKa = 3.2 pKa = -7 pKa = -9 pKa = -10

Acidity increase

The important factor is the strength of the H-X bond, the stronger the bond, the weaker the acid.

Because HI, HBr, and HCl are such stronger acids, their conjugate

bases (I-, Br-, Cl-) are all very weaker bases.

Basicity increase

F- Cl- Br- I-

We see the same trend of acidities and basicities in other vertical columns of the periodic table. For example:

Acidity increase

Basicity increase

H2O H2S H2Se

OH- SH- SeH-

Lets see an example of CH4, NH3, H2O and HF. These compoundsare all hydrides of first-row elements and electronegaticity increasesacross a row of the periodic table from left to right.

Electronegaticity increase

C N O F

Because fluorine is the most electronegative, the bond in H-F is Most polarized, and the proton in H-F is the most positive, so H-F is the most acidic.

When we compare compounds in the same horizontal row of the periodic table, the dominant factor becomes the electronegativity of the atom bond to the hydrogen.

CH3—H NH2—H HO—H F—HpKa = 48 pKa = 38 pKa = 15.7 pKa = 3.2

Acidity increase

Because H-F is the strongest acid, its conjugate base, the fluorideion will be the weakest base, fluorine is the most electronegative atom and it accommodates the negative charge most readily.

Basicity insrease

CH3- NH2

- HO- F-

3.5A THE EFFECT OF HYBRIDIZATION

The protons of ethyne are more acidic than those of ethene, whichin turn, are more acidic than of ethane.

H C C H C C

H

H H

H

C CH

HH

H

H

HEthynepKa = 25

EthenepKa = 44

EthanepKa = 50

We can explain this order of acidities on the basis of the hybridiza-tion state of carbon in each compound. Electrons of 2s orbitals have Lower energy than those of 2p orbitals because in 2s orbitals tend to be much closer to the nucleus than electrons in 2p orbitals.

With hydrid, therefore, having more s charater means that the electrons of the anion will, on the average, be lower in enegy,andthe anion will be more state.

Now we can see how the order of relative acidities of ethyne, ethene,and ethane parallels the effective electronegativity of the carbon atom in each compound:

Relative acidity:

HC CH H2C CH2 H3C CH3£¾ £¾

The more the electronegative, the most positive the hydrogens. So ethyne donates a proton to a base more readily. And in the sameway, the ethynide ion id the weaker base because the more electro-negative carbon of ethyne is best able to stabilize the negativecharge.

Relative basicity:

HC C-H2C CH-H3C CH2- £¾ £¾

3.5B INDUCTIVE EFFECTS

The carbon-carbon bond of ethane is completely nonpolar becauseat each end of the bond there are two equivalent methyl groups.

H3C CH3

Ethyne

The C-C bond is nonpolar

This is not the case with ethyl fluoride, however:

£¾ £¾CH2 FH3C

2 1

Inductive effect: the effect of the polarization of the carbon-carbonbond results from an intrinsic electron-at-tracting ability of the fluorine that is transmitted through space and through the bonds of the molecule.

Inductive effect weaken steadily as the distance from the substituentincrease. In this instance, the positive charge that the fluorine impartsto C-1 is greater than that imparted to C-2 because the fluorine is closer to C-1.

Inductive effects help us to understand why carboxylic acids are much more acidic than alcohols.

H3C C OH

O

H3CH2C OH

Acetivc acid pKa = 4.74

Ethyl alcohol pKa = 15.9

The key to the much greater acidity of acetic acid is the powerelectro-attracting inductive effect of its carbonyl group(C=O) whencompare to the CH2 group in the corresponding position of ethylalcohol.

The carbonyl group of acetic acid, because its bears a large positivecharge, adds its electron-attracting effect to that of the oxygen of the hydroxyl group attached to it; this makes the hydroxyl proton much more positive than the proton of the alcohol.

This greater positive charge on the proton of the acid means that theproton separates more readily. The electron-attracting effect of the carbonyl group also stabilizes the anion that forms from the carboxylic acid, and, therefore, the carboxylate ion is weaker base than the ethoxide ion.

H3C C O

O

H3CH2C OH H

H3C C O

O

H3CH2C O£¼

£¼ £¼

£¼£¼

£¼

£¼

Stronger acid Weaker acid

Weaker base Stronger base

The acid-strengthen effect of electron-attracting group can also be shown by comparing the acidities of acetic acid and chloroacetic acid.

H3C C OH

O

£¼£¼

£¼

£¼H2C C OH

O

£¼

£¼

Cl

pKa = 4.76 pKa = 2.86

The greater acidity of chloroacetic acid can be attributed, in part, to the extra electron-attracting inductive effect of the electronegative chlorine atom.

So, dispersal of charge always makes a species more stable, andany factor that stabilizes the conjugate base of an acid will increasethe strength of the acid.

3.6 THE RELATIONSHIP BETWEEN THE EQUILIBRIUM CONSTANT AND THE STANDARD FREE-ENERGY CHANGE, △ Gº

An important relationship exists between the equilibrium constant and the standard free-energy change that accompanies the reaction.

△ Gº = -2.303RTlogKeq

R is the gas constant and equals 1.987 cal /(K·mol); T is the absolute temperature in kelvins (K).

A negative value of △ Gº is associated with reactions that favor the formation of products when equilibrium is reached, and a positivevalue of △ Gº is associated with reactions for which the formationof the products at equilibrium is unfavorable.

The free-energy change(△ Gº) has two components, the enthalpy change(△ Hº) and the entropy chang(△ Sº). The relationship between these thermodynamic quantities is:

△ Gº = △ Hº - T△ Sº

△ Hº is associated with changes in bonding that occur in a reaction.If, collectively, stronger bonds are formed in the products than exitedin the starting materials, then △ Hº will be negative,vice versa. A negative value for △ Hº , therefore, will contribute to making △ Gº negative, and will, consequently favor the formation of products.

The more random a system is, the greater is its entropy. A positiveentropy change (+ △ Sº ) is always associated with a change from a more ordered system to a cess. We can see from the aboverelationship that a positive entropy change makes a negative contri-bution to △ Gº and is energetically favorable for the formation ofproducts.

3.7 THE EFFECT OF THE SOLVENT ON A △ Hº CIDITY

In the absence of a solvent, most acids are far weaker than they are in solution. For example: in gas phase acetic acid is estimated to have a pKa of about 130! The reason: when an acetic acid moleculedonates a proton to a water molecule in the gas phase, the ions that are formed are oppositely charged particles and these particles mustbecome separated.

H3C C

O

OH + H2O H3C C

O

O- + H3O+

A protic solvent: solvent that has a hydrogen atom attached to astrongly electronegative element such an oxygen or nitrogen.

Molecules of a protic solvent can form hydrogen bonds to the un-shared electron pairs of oxygen atoms of an acid and its conjugate base, but they may not stabilize both equally. For example: if aceticacid in aqueous solution, hydrogen bonding to CH3COO¯ is much stronger than to CH3COOH because the water molecules are more attracted by the negative charge.

This differential solvation, moreover, has important consequences forthe entropy charge that accompanies the ionization. Solvation of any species decreases the entropy of the solvent because the solvent mole-cules become much more ordered as they surround molecules of the solute.

The following table list the thermodynamic values for the dissociationof acetic and chloroacetic acids in H2O at 25℃

___________________________________________________________________________________

___________________________________________________________________________________

___________________________________________________________________________________

ACID

CH3CO2H

ClCH2CO2H

pKa ¡÷Go ¡÷Ho T¡÷So(kcal mol-1) (kcal mol-1) (kcal mol-1)

4.76

2.86

+ 6.5

+3.9

- 0.1

- 1.1

+ 6.6

+ 5.0

3.8 ACID AND BASES IN NONAQUEOUS AOLUTIONS

The reaction of amide ion in aqueous solution as the following:

H O H NH2- H O

- NH3++

Stronger acid pKa = 15.7

Stronger base Weaker base Weaker acid pKa = 38

This example illustrate what is called the leveling effect of the solvent. The solvent here, water, converts any base stronger thana hydroxide ion to a hydroxide ion by donating.

But we can convert ethyne to its conjugate base by treating it withsodium amide in liquid ammonia.

NH2-

NH3++

Stronger acid pKa = 25

Stronger base(from NaNH2)

Weaker base

Weaker acid pKa = 38

H C C H H C C-liquidNH3

Most alkynes with a proton attached to a triply bonded carbonhave pKa values of about 25, therefore, all react with sodiumamide in liquid ammonia in the same way that ethyne dose:

NH2-

NH3++

Stronger acid pKa = 25

Stronger base

Weaker base

Weaker acid pKa = 38

R C C H R C C-liquidNH3

Alcohols are often used as solvents for organic reactions because being somewhat less polar than water, they dissolve less polar organic compound.

-+ +

Stronger acid pKa = 16

Stronger base (from NaH)

Weaker base

Weaker acid pKa = 35

CH3CH2O H- ethyl alcohol CH3CH2O H2H

Another example:

-+ +

Stronger acid pKa = 17

Stronger base (from NaH)

Weaker base

Weaker acid pKa = 35

(CH3)3CO H H- ethyl alcohol (CH3)3CO H2

Alkyllithiums react as though they contained alkanide ions, and being the conjugate bases of alkanes, alkanide ions are the strongestbases that we shall encounter.

-+ +

Stronger acid pKa = 25 (from CH3CH2Li)

Weaker base

Weaker acid pKa = 50

R C C H R C C-CH2CH3hexane CH3CH3

Stronger base

Alkyllithiums can be easily prepared by allowing an alkyl bromide to react with lithium metal in an ether solvent.

RBr + 2Li RLi + LiBrdiethyl ether

General reaction:

Specific reaction:

CH3CH2Br + 2Li CH3CH2Li + LiBrdiethyl ether

3.9 ACID-BASE REACTIONS AND THE SYNTHESIS OF DEUTERIUM AND TRITIUM-LABELED COMPOUNDS

Chemists often use compounds in which deuterium(²H) or tritium(³H)atoms have replaced one or more hydrogen atoms of the compound as a method of “labeling” or identifying particular hydrogen atoms.

The extra mass and additional neutrons associated with a deuteriumor tritium atom often makes its position in a molecule easy to locate by certain spectroscopic.

One way to introduce a deuterium or tritium atom into a specific location in a molecule is through the acid-base reaction that takes place when a very strong base is treated with D2O or T2O. For example:

Stronger acid

Weaker base

hexaneCH3CH2

CH3

Li+- + D2O CH3CH

CH3

D + OD-

Isopropyllithium£¨ Òì±û»ùﮣ© stronger base

2-deuteriopeopane (¶þë®±ûÍé £© Weaker acid

3.10 SOME IMPORTANT TERMS AND CONCEPTS

A Brnsted-Loery acid : a substance that donate a proton; A Brnsted-Lowry base: a substance that can accept a proton.A Lewis acid: an electron-pair acceptor, a Lewis base is an electron-pairdoner.Curved arrows: used to show the direction of electron flow when mechanisms are written.The strength of an acid can be expressed by its acidity constant, Ka.

Ka =[HA]

[H3O+] [A-]

Or by its pKa: pKa = -logKa

The strength of a base: inversely related to the strength of its conju- gated acid; the weaker the conjugated acid, the stronger the base.The outcome of the acid-base reaction:predicted on the basis of the principle that acid-base reactions proceed toward equilibrium so as to favor formation of the weaker acid and the weaker base.An inductive effect:reflects the ability of a substituent to attract or release electrons because of its electronegativity.

Dispersal of electrical charge: always makes a chemical entity more stable.The relationship between Keq and the standard free-energy change(△ Gº ):

△ Gº = -2.303RTlogKeq

The relationship of △ Gº , △ Hº and △ Sº :

△ Gº = △ Hº - T△ Sº

A protic solvent: one that has a hydrogen atom attached to a strongly electronegative atom.