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BrØnsted-Lowry Acids and Bases• A BrØnsted-Lowry acid is a proton donor.

• A BrØnsted-Lowry base is a proton acceptor.

• H+ = proton

Acids and Bases

BrØnsted-Lowry Acids and Bases

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.

Acids and Bases

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Reactions of BrØnsted-Lowry Acids and Bases• A BrØnsted-Lowry acid base reaction results in the transfer of

a 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 H—A loses a proton, leaving the electron pair in theH—A bond on A.

Acids and Bases

Reactions of BrØnsted-Lowry Acids and Bases• The movement of electrons in reactions can be illustrated using

curved arrow notation. Because two electron pairs are involved inthis 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 andproducts to indicate that the reaction can proceed in the forwardand reverse directions. These are equilibrium arrows.

Examples:

Acids and Bases

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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, the stronger an acidit is.• Acidity is measured by an equilibrium constant

• When a BrØnsted-Lowry acid H—A is dissolved in water, an acid-base reaction occurs, and an equilibrium constant can be written forthe reaction.

Acids and Bases

Acid Strength and pKa

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 use“pKa” values than Ka values.

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Acid Strength and pKa

Acids and Bases

Factors that Determine Acid Strength Anything that stabilizes a conjugate base A:¯ makes the

starting acid H—A more acidic.

Four factors affect the acidity of H—A. These are:

o Element effects

o Inductive effects

o Resonance effects

o Hybridization effects

No matter which factor is discussed, the same procedure isalways 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 theacid.

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Factors that Determine Acid Strength

Element Effects—Trends in the Periodic Table.

Across a row of the periodic table, the acidity of H—Aincreases as the electronegativity of A increases.

Positive or negative charge is stabilized when it is spreadover a larger volume.

Acids and Bases

Factors that Determine Acid StrengthElement Effects—Trends in the Periodic Table.

• Down a column of the periodic table, the acidity of H—Aincreases as the size of A increases.

• Size, and not electronegativity, determines acidity down acolumn.

• The acidity of H—A increases both left-to-right across a rowand down a column of the periodic table.

• Although four factors determine the overall acidity of aparticular hydrogen atom, element effects—the identity ofA—is the single most important factor in determining theacidity of the H—A bond.

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Factors that Determine Acid Strength—Inductive Effects

• An inductive effect is the pull of electron densitythrough σ bonds caused by electronegativitydifferences in 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.

Acids and Bases

Factors that Determine Acid Strength—Inductive Effects

Factors that Determine Acid Strength—Inductive Effects• When electron density is pulled away from the negative

charge through σ bonds by very electronegative atoms, it isreferred to as an electron withdrawing inductive effect.

• More electronegative atoms stabilize regions of high electrondensity 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 H—A increases with the presence of electronwithdrawing groups in A.

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Factors that Determine Acid Strength—Resonance 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 enjoysresonance stabilization, whereas that of ethanol does not.

Acids and Bases

• Resonance delocalization makes CH3COO¯ more stablethan CH3CH2O¯, so CH3COOH is a stronger acid thanCH3CH2OH.

• The acidity of H—A increases when the conjugatebase A:¯ is resonance stabilized.

Factors that Determine Acid Strength—ResonanceEffects

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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¯.

Factors that Determine Acid Strength—ResonanceEffects

Acids and Bases

Factors that Determine Acid Strength—HybridizationEffects

• The final factor affecting the acidity of H—A 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 C—H bonds.

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Factors that Determine Acid Strength—HybridizationEffects

Acids and Bases

Acids and Bases

Factors that Determine Acid Strength—HybridizationEffects

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Commonly Used Acids in Organic ChemistryIn addition to the familiar acids HCl, H2SO4 and HNO3, anumber of other acids are often used on organicreactions. Two examples are acetic acid and p-toluene-sulfonic acid (TsOH).

Acids and Bases

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

Commonly Used Acids 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.

Commonly Used Acids in Organic Chemistry

Acids and Bases

Lewis Acids and Bases• The Lewis definition of acids and bases is more general

than the BrØnsted-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 BrØnsted-Lowrybases. Both have an available electron pair—a lone pair oran electron pair in a π bond.

• A BrØnsted -Lowry base always donates this electron pairto a proton, but a Lewis base donates this electron pair toanything that is electron deficient.

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Lewis Acids and Bases• A Lewis acid must be able to accept an electron pair, but

there are many ways for this to occur.

• All BrØnsted-Lowry acids are also Lewis acids, but thereverse 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 BrØnsted-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.

Acids and Bases

Lewis Acids and Bases• Any reaction in which one species donates an electron pair

to 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 ofBF3 with H2O. H2O donates an electron pair to BF3 to form anew bond.

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Lewis Acids and Bases• A Lewis acid is also called an electrophile.

• When a Lewis base reacts with an electrophile otherthan a proton, the Lewis base is also called anucleophile. In this example, BF3 is the electrophile andH2O is the nucleophile.

electrophile nucleophile

Acids and Bases

• 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.

Lewis Acids and Bases

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• In some Lewis acid-base reactions, one bond is formedand one bond is broken. To draw the products of thesereactions, keep in mind the following steps:

o Always identify the Lewis acid and base first.

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

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

Lewis Acids and Bases

Acids and Bases

Lewis Acids and Bases

Consider the Lewis acid-base reaction betweencyclohexene and H—Cl. The BrØnsted-Lowry acid HClis also a Lewis acid, and cyclohexene, having a π bond,is the Lewis base.

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Lewis Acids and Bases• To draw the product of this reaction, the electron pair in

the π bond of the Lewis base forms a new bond to theproton of the Lewis acid, generating a carbocation.

• The H—Cl bond must break, giving its two electrons toCl, forming Cl¯.

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

Acids and Bases