Structural and Electronic Effects Acids and Bases

Lewis acid – a substance that accepts an electron pair • All BrOnsted-Lowry acids are also Lewis acids, but the reverse is not necessarily true. Any species that is electron deficient and capable of accepting an electron pair is also a Lewis acid. • Common examples of Lewis acids (which are not BrOnstedLowry acids) include BF3 and AlCl3. These compounds contain elements in group 3A of the periodic table that can accept an electron pair because they do not have filled valence shells of electrons. • Lewis base – a substance that donates an electron pair 3 4 • In a Lewis acid-base reaction, a Lewis base donates an electron pair to a Lewis acid. • One bond is formed and no bonds are broken.This is illustrated in the reaction of BF3 with H2O. H2O donates an electron pair to BF3 to form a new bond. pKa’s AND ACID STRENGTH pKa – An alternative to Ka to describe acid strength. (A concise way to state the strength of an acid. ) electron pair is not removed from the Lewis base. Instead, it is donated to an atom of the Lewis acid and one new covalent bond is formed. 5 1 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 acid it is. • Acidity is measured by an equilibrium constant. When a Bronsted-Lowry acid H—A is dissolved in water, an acid-base reaction occurs, and an equilibrium constant can be written for the reaction. Because the concentration of the solvent H2O is essentially constant, the equation can be rearranged and a new equilibrium constant, called the acidity constant, Ka, can be defined. It is generally more convenient when describing acid strength to use “pKa” values than Ka values. 7 8 COMPARISON OF pKa and Ka VALUES pKa = – log Ka strong acids weak acids pKa -2 0 2 10-2 4 6 10-6 8 10 10-10 12 14 10-14 2 Ka 10 The smaller the value of the pKa the stronger the acid.We will use pKa to describe the strengths of acids. 10 EVALUATION OF ACID STRENGTH Commonly Used Bases in Organic Chemistry Common strong bases used in organic reactions are more varied in structure. HA + H2O H3O+ + A- In water, all acids form hydronium ion, the important factor of difference is the conjugate base. The difference between a strong acid and a weak acid is in the stability of the conjugate base. AA- E N E R G Y 11 WEAK ACID has strong conj. base (=higher energy) STRONG ACID has weak conj. base (=lower energy) HA ionization easier 2 EVALUATION OF ACID STRENGTH The weaker an acid, the stronger is its conjugate base; • The stronger an acid, the weaker is its conjugate base. • HCl is a very strong acid; it gives up its proton readily; its conjugate base, Cl-, has very little affinity for H+. It is a stable CB 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: Element effects Inductive effects Resonance 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. 14 • CH3CO2H is a moderately weak acid; it gives up its proton somewhat reluctantly; its conjugate base, CH3CO2-, is weakly basic and has a modest affinity for H+. CB is relatively unstable 13 Element Effects—Trends in the Periodic Table. Factors that Determine Acid Strength: Across a row of the periodic table, the acidity of H—A increases as the electronegativity of A increases. Acids and BasesFactors that Determine Acid Strength—Inductive Effects • An inductive effect is the pull of electron density through ? bonds caused by electronegativity differ-ences between atoms. • In the example below, when we compare the acidities of ethanol and 2,2,2-trifluoroethanol, we note that the latter is more acidic than the former. Positive or negative charge is stabilized when it is spread over a larger volume. 15 16 • When electron density is pulled away from the negative charge 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 the site of the negative charge, the greater the effect. • The acidity of H—A increases with the presence of electron withdrawing groups in A. TYPES OF INDUCTIVE EFFECTS ELECTRON WITHDRAWING GROUPS ELECTRON DONATING GROUPS ?Cl ?+ C ?+ CH3 ?C F, Cl, Br, N, O electronegative elements take electron density from cabon 17 R, CH3, B, Si alkyl groups and elements less electronegative than carbon donate electron density to carbon INDUCTIVE EFFECTS HALOACIDS Chlorine helps to stabilize -CO2by withdrawing electrons Inductive effect in acids and bases O H H C H O C O H F H C H O C O H H3C H C H O C O H ?- ? + C Cl O C O O ?+ ?- ? + CH3-CO2 FCH2-CO2 CH3CH2-CO2CH3, electron donating ?Cl ?+ C ?- No substituent F, electron withdrawing C F makes O-H bond weak, easily broken – more acidic CH3 makes bond stronger, difficult to break – less acidic Strength of acids can also be determined by the stability of A:20 The effect diminishes with distance – it carries for about 3 bonds.INDUCTIVE EFFECTS – 1 (EWG) increasing electronegativity INDUCTIVE EFFECTS – 2 (EDG) pKa Values H COOH CH3 CH3CH2 CH3CH2CH2 CH3 COOH COOH COOH pKa Values 3. 13 2. 87 2. 81 2. 66 CH3 COOH Cl CH2 COOH Cl Cl CH Cl Cl C COOH Cl COOH multiple substituents I Br Cl F CH2COOH CH2COOH CH2COOH CH2COOH 4. 75 2. 81 1. 29 0. 65 3. 75 4. 75 4. 87 4. 81 5. 02 CH3CH2CH2 COOH CH2CH2CH2 COOH Cl CH3 CH CH2 Cl CH3CH2 CH COOH Cl COOH 4. 8 4. 5 4. 0 2. 9 distance CH3 C COOH CH3 ELECTRONELECTRON-WITHDRAWING EFFECTS STRENGTHEN ACIDS O (-) R C OC OO (-) O OR S OOConversely ….. ELECTRONELECTRON-DONATING EFFECTS WEAKEN ACIDS O (-) R C OC OO (-) OOR S OO – – Any effect that “bleeds” electron density away from the negatively-charged end of the conjugate base will stabilize (lower the energy) of the conjugate base and therefore make the acid stronger. Any effect that “pushes” extra electron density toward the negatively-charged end of the conjugate base will destabilize (increase the energy) of the conjugate base and make the acid weaker. 4 Example 1: Compare the acids of H2O, HOCl and CH3OH Dissociation productsExample 2: Compare the acids of CH3OH, and (CH3)3COH Dissociation products H 2O HOCl CH3OH H OO O- H+ + OHH+ + ClOH+ + CH3O- CH3OH (CH3)3COH CH3 H+ + CH3OH+ + C(CH3)3O- NO EWG or EDG CH3 is EDG, therefore anion is unstable Cl is EWG, therefore anion is more stable Compare stability Compare stability CH3 Cl – CH3 O- H3C C CH3 O- Both have EDGs which leads to anion instability 1 x CH3 – 3 x CH3, ClO > OH > CH3O most stable to least stable anion – HOCl > H2O > CH3OH acid strengths 25 Therefore (CH3)3CO anion is less stable than CH3O- anion. (CH3)3COH weaker acid 26 5

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