Chapter 2 Lecture_2

8/11/2019 Chapter 2 Lecture_2

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Brnsted-Lowry Acids and Bases

A Brnsted-Lowry acid is a proton donor.

A Brnsted-Lowry base is a proton acceptor.

H+= proton

Acids and Bases

Figure 2.1Examples of BrnstedLowry

acids and bases


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Some molecules contain both hydrogen atoms and lonepairs and thus, can act either as acids or bases,depending on the particular reaction. An example is theaddictive pain reliever morphine.


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Which of the following molecules is not a Bromsted-Lowry acid?

HBr NH3 CCl4

Bromsted-Lowry acid mustcontain a proton.

Which of the following is not a Bromsted-Lowry base?

H3C CH3 O

Bromsted-Lowry bases must have a lone pair or pibond to donate.


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Reactions of Brnsted-Lowry Acids and Bases

A Brnsted-Lowry acid base reaction results in the transfer ofa proton from an acid to a base.

In an acid-base reaction, one bond is broken, and anotherone is formed.

The electron pair of the base B: forms a new bond to theproton of the acid.

The acid HA loses a proton, leaving the electron pair in theHA bond on A.


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The movement of electrons in reactions can be illustrated usingcurved arrow notation. Because two electron pairs are involved

in this reaction, two curved arrows are needed. Loss of a proton from an acid forms its conjugate base.

Gain of a proton by a base forms its conjugate acid.

A double reaction arrow is used between starting materials and

products to indicate that the reaction can proceed in the forwardand reverse directions. These are equilibrium arrows.

Examples:


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Label the acid, base, conjugate acid and conjugatebase in the following reaction.

O

H3C CH3

+ O CH3

O

H3C CH2

+HO CH3

A BCB CA

O

H3C CH2

+ O CH3

O

H3C CH2

+HO CH3

H

Use curved arrows to show the movement of electronpairs.


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Draw the products of the following reaction ( use curvedarrows to show movement of electron pairs)

H3C NH2 + HCl

H3C NH2 + H Cl H3C NH3

+ Cl

H H + H

H H + H H

+ H H


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Acid Strength and pKa

Acid strength is the tendency of an acid to donate a proton.

The more readily a compound donates a proton, thestronger an acid it is.

Acidity is measured by an equilibrium constant.

When a Brnsted-Lowry acid HA is dissolved in water, an

acid-base reaction occurs, and an equilibrium constant canbe written for the reaction.


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Because the concentration of the solvent H2O is essentially constant,the equation can be rearranged and a new equilibrium constant, calledthe acidity constant, Ka, can be defined.

It is generally more convenient when describing acid strength to usepKavalues than Kavalues.


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Factors that Determine Acid Strength

Anything that stabilizes a conjugate base A: makes thestarting acid HA more acidic.

Four factors affect the acidity of HA. These are:

Element effects

Inductive effectsResonance effects

Hybridization effects

No matter which factor is discussed, the same procedure is

always followed. To compare the acidity of any two acids:o Always draw the conjugate bases.

o Determine which conjugate base is more stable.

o The more stable the conjugate base, the more acidic the

acid.


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Element EffectsTrends in the Periodic Table.

Across a row of the periodic table, the acidity of HAincreases as the electronegativity of A increases.

Positive or negative charge is stabilized when it is spread

over a larger volume.


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Down a column of the periodic table, the acidity of HAincreases as the size of A increases.

Size, and not electronegativity, determines acidity down acolumn.

The acidity of HA increases both left-to-right across a row

and down a column of the periodic table. Although four factors determine the overall acidity of a

particular hydrogen atom, element effectsthe identity ofAis the single most important factor in determining theacidity of the HA bond.


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Factors that Determine Acid StrengthInductive Effects

An inductive effect is the pull of electron densitythrough

bonds caused by electronegativity differ-ences between atoms.

In the example below, when we compare the aciditiesof ethanol and 2,2,2-trifluoroethanol, we note that thelatter is more acidic than the former.


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The reason for the increased acidity of 2,2,2-trifluoroethanol is that the three electronegativefluorine atoms stabilize the negatively chargedconjugate base.


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When electron density is pulled away from the negativecharge through bonds by very electronegative atoms, it is

referred to as an electron withdrawing inductive effect. More electronegative atoms stabilize regions of high electron

density by an electron withdrawing inductive effect.

The more electronegative the atom and the closer it is to thesite of the negative charge, the greater the effect.

The acidity of HA increases with the presence of electronwithdrawing groups in A.


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Factors that Determine Acid StrengthResonance Effects

Resonance is a third factor that influences acidity.

In the example below, when we compare the acidities ofethanol and acetic acid, we note that the latter is more acidicthan the former.

When the conjugate bases of the two species are compared, itis evident that the conjugate base of acetic acid enjoys

resonance stabilization, whereas that of ethanol does not.


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Resonance delocalization makes CH3COO more

stable than CH3CH2O, so CH3COOH is a stronger acid

than CH3CH2OH.

The acidity of HA increases when the conjugatebase A:is resonance stabilized.


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Electrostatic potential plots of CH3CH2O and

CH3COO below indicate that the negative charge isconcentrated on a single O in CH3CH2O, butdelocalized over both of the O atoms in CH3COO.


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Factors that Determine Acid StrengthHybridizationEffects

The final factor affecting the acidity of HA is thehybridization of A.

The higher the percent of s-character of the hybridorbital, the closer the lone pair is held to the nucleus,and the more stable the conjugate base.

Let us consider the relative acidities of three differentcompounds containing CH bonds.


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Figure 2.5Summary of the factors that

determine acidity


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Calculate the pkas and which is the strongeracid?

OH CH3

ka= 10-10 10-41

pka= 10 41

OH


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Which is the strongest base?

H20 (pka= 15.7) NH3(pka= 38)

CH4(pka= 58)CH4

Draw the products and determine the direction of the

equilibrium.

CH4 + OH CH3

+H2O

pka= 50 15.7


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Which of these bases is strong enough to deprotonateCH3COOH (pka= 4.8) ?

H Cl

pka= 25 35 -7

Both acetylene and hydrogen will work(must have higher pk

a

)

Not using pka, which is the stronger acid?

H2S HBr

HBr, b/c Br is further to the right and further down inthe periodic table.


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Which is the stronger acid?

F2CHCH2COOH OR Cl2CH2COOH

F2CHCH2COOH is the most acidic b/c acids arestabilized by electron withdrawing groups and F is moreelectronegative than Cl

Which is the stronger acid and why?

H

H

H

B/c the C-H bond involves sp2hybridized orbital instead ofsp3


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Commonly Used Acids in Organic Chemistry

In addition to the familiar acids HCl, H2SO4and HNO3, anumber of other acids are often used in organicreactions. Two examples are acetic acid and p-toluene-sulfonic acid (TsOH).


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Common strong bases used in organic reactions aremore varied in structure.

Commonly Used Bases in Organic Chemistry


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It should be noted that:

Strong bases have weak conjugate acids with high pKavalues, usually > 12.

Strong bases have a net negative charge, but not allnegatively charged species are strong bases. For example,none of the halides F, Cl, Br, or I, is a strong base.

Carbanions, negatively charged carbon atoms, are especiallystrong bases. A common example is butyllithium.

Two other weaker organic bases are triethylamine andpyridine.


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Lewis Acids and Bases

The Lewis definition of acids and bases is more general

than the Brnsted-Lowry definition. A Lewis acid is an electron pair acceptor.

A Lewis base is an electron pair donor.

Lewis bases are structurally the same as Brnsted-Lowry

bases. Both have an available electron paira lone pair oran electron pair in a bond.

A Brnsted-Lowry base always donates this electron pairto a proton, but a Lewis base donates this electron pair to

anything that is electron deficient.


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A Lewis acid must be able to accept an electron pair, butthere are many ways for this to occur.

All Brnsted-Lowry acids are also Lewis acids, but the

reverse is not necessarily true.

o Any species that is electron deficient and capable ofaccepting an electron pair is also a Lewis acid.

Common examples of Lewis acids (which are not Brnsted-

Lowry acids) include BF3 and AlCl3. These compoundscontain elements in group 3A of the periodic table that canaccept an electron pair because they do not have filledvalence shells of electrons.


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Any reaction in which one species donates an electron pairto another species is a Lewis acid-base reaction.

In a Lewis acid-base reaction, a Lewis base donates anelectron pair to a Lewis acid.

Lewis acid-base reactions illustrate a general pattern inorganic chemistry. Electron-rich species react with electron-poor species.

In the simplest Lewis acid-base reaction one bond is formedand no bonds are broken.This is illustrated in the reaction ofBF3with H2O. H2O donates an electron pair to BF3 to form anew bond.


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Two other examples are shown below. Note that in eachreaction, the electron pair is not removed from theLewis base. Instead, it is donated to an atom of theLewis acid and one new covalent bond is formed.


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In some Lewis acid-base reactions, one bond is formedand one bond is broken. To draw the products of these

reactions, keep in mind the following steps:Always identify the Lewis acid and base first.

Draw a curved arrow from the electron pair of thebase to the electron-deficient atom of the acid.

Count electron pairs and break a bond when neededto keep the correct number of valence electrons.


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To draw the product of this reaction, the electron pair inthe bond of the Lewis base forms a new bond to the

proton of the Lewis acid, generating a carbocation.

The HCl bond must break, giving its two electrons toCl, forming Cl.

Because two electron pairs are involved, two curvedarrows are needed.


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Which is not a Lewis base?

NH3

Lewis base must have an electron pair to donate.

Which is not a Lewis acid?

Br

B

Br Br

OH CH

None, all are Lewis acids b/c they either contain anunfilled valence shell or an acidic proton.


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Chapter 2 Lecture_2

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