1. 1 Organic Chemistry, Second Edition Janice Gorzynski Smith University of Hawai’i Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Chapter 2 Lecture Outline Prepared by Rabi Ann Musah State University of New York at Albany

2. 2 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 Figure 2.1 Examples of Brønsted–Lowry acids and bases

3. 3 Brønsted-Lowry Acids and Bases Some molecules contain both hydrogen atoms and lone pairs and thus, can act either as acids or bases, depending on the particular reaction. An example is the addictive pain reliever morphine. Acids and Bases

4. 4 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 another one is formed. The electron pair of the base B: forms a new bond to the proton of the acid. The acid H—A loses a proton, leaving the electron pair in the H—A bond on A. Acids and Bases

5. 5 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 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 forward and reverse directions. These are equilibrium arrows. Examples: Acids and Bases

6. 6 Acid Strength and pK a 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 Brønsted-Lowry acid H—A is dissolved in water, an acid-base reaction occurs, and an equilibrium constant can be written for the reaction. Acids and Bases

7. 7 Acid Strength and pK a Because the concentration of the solvent H 2 O is essentially constant, the equation can be rearranged and a new equilibrium constant, called the acidity constant, K a, can be defined. It is generally more convenient when describing acid strength to use “pK a ” values than K a values. Acids and Bases

8. 8 Acid Strength and pK a Acids and Bases

9. 9 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: oAlways draw the conjugate bases. oDetermine which conjugate base is more stable. oThe more stable the conjugate base, the more acidic the acid. Acids and Bases

10. 10 Factors that Determine Acid Strength Element Effects—Trends in the Periodic Table. Across a row of the periodic table, the acidity of H—A increases as the electronegativity of A increases. Positive or negative charge is stabilized when it is spread over a larger volume. Acids and Bases

11. 11 Factors that Determine Acid Strength Element Effects—Trends in the Periodic Table. Down a column of the periodic table, the acidity of H—A increases as the size of A increases. Size, and not electronegativity, determines acidity down a column. The acidity of H—A 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 effects—the identity of A—is the single most important factor in determining the acidity of the H—A bond. Acids and Bases

12. 12 Factors 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. Acids and Bases

13. 13 The reason for the increased acidity of 2,2,2- trifluoroethanol is that the three electronegative fluorine atoms stabilize the negatively charged conjugate base. Acids and Bases Factors that Determine Acid Strength—Inductive Effects

14. 14 Factors that Determine Acid Strength—Inductive Effects 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. Acids and Bases

15. 15 Factors that Determine Acid Strength—Resonance Effects Resonance is a third factor that influences acidity. In the example below, when we compare the acidities of ethanol and acetic acid, we note that the latter is more acidic than the former. When the conjugate bases of the two species are compared, it is evident that the conjugate base of acetic acid enjoys resonance stabilization, whereas that of ethanol does not. Acids and Bases

16. 16 Resonance delocalization makes CH 3 COO ¯ more stable than CH 3 CH 2 O ¯, so CH 3 COOH is a stronger acid than CH 3 CH 2 OH. The acidity of H—A increases when the conjugate base A: ¯ is resonance stabilized. Factors that Determine Acid Strength—Resonance Effects Acids and Bases

17. 17 Electrostatic potential plots of CH 3 CH 2 O ¯ and CH 3 COO ¯ below indicate that the negative charge is concentrated on a single O in CH 3 CH 2 O ¯, but delocalized over both of the O atoms in CH 3 COO ¯. Factors that Determine Acid Strength—Resonance Effects Acids and Bases

18. 18 Factors that Determine Acid Strength—Hybridization Effects The final factor affecting the acidity of H—A is the hybridization of A. The higher the percent of s-character of the hybrid orbital, 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 different compounds containing C—H bonds. Acids and Bases

19. 19 Factors that Determine Acid Strength—Hybridization Effects Acids and Bases

20. 20 Acids and Bases Factors that Determine Acid Strength—Hybridization Effects Figure 2.5 Summary of the factors that determine acidity

21. 21 Commonly Used Acids in Organic Chemistry In addition to the familiar acids HCl, H 2 SO 4 and HNO 3, a number of other acids are often used in organic reactions. Two examples are acetic acid and p-toluene- sulfonic acid (TsOH). Acids and Bases

22. 22 Common strong bases used in organic reactions are more varied in structure. Commonly Used Bases in Organic Chemistry Acids and Bases

23. 23 It should be noted that: Strong bases have weak conjugate acids with high pK a values, usually > 12. Strong bases have a net negative charge, but not all negatively 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 especially strong bases. A common example is butyllithium. Two other weaker organic bases are triethylamine and pyridine. Commonly Used Bases in Organic Chemistry Acids and Bases

24. 24 Lewis Acids and Bases The Lewis definition of acids and bases is more general than the Br Ø nsted-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 Br Ø nsted-Lowry bases. Both have an available electron pair—a lone pair or an electron pair in a  bond. A Br Ø nsted-Lowry base always donates this electron pair to a proton, but a Lewis base donates this electron pair to anything that is electron deficient. Acids and Bases

25. 25 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 Br Ø nsted-Lowry acids are also Lewis acids, but the reverse is not necessarily true. oAny 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 Br Ø nsted- Lowry acids) include BF 3 and AlCl 3. 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. Acids and Bases

26. 26 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 an electron pair to a Lewis acid. Lewis acid-base reactions illustrate a general pattern in organic chemistry. Electron-rich species react with electron- poor species. In the simplest Lewis acid-base reaction one bond is formed and no bonds are broken. This is illustrated in the reaction of BF 3 with H 2 O. H 2 O donates an electron pair to BF 3 to form a new bond. Acids and Bases

27. 27 Lewis Acids and Bases A Lewis acid is also called an electrophile. When a Lewis base reacts with an electrophile other than a proton, the Lewis base is also called a nucleophile. In this example, BF 3 is the electrophile and H 2 O is the nucleophile. electrophilenucleophile Acids and Bases

28. 28 Two other examples are shown below. Note that in each reaction, the 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. Lewis Acids and Bases Acids and Bases

29. 29 In some Lewis acid-base reactions, one bond is formed and 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 the base to the electron-deficient atom of the acid. Count electron pairs and break a bond when needed to keep the correct number of valence electrons. Lewis Acids and Bases Acids and Bases

30. 30 Lewis Acids and Bases Consider the Lewis acid-base reaction between cyclohexene and H—Cl. The Br Ø nsted-Lowry acid HCl is also a Lewis acid, and cyclohexene, having a  bond, is the Lewis base. Acids and Bases

31. 31 Lewis Acids and Bases To draw the product of this reaction, the electron pair in the  bond of the Lewis base forms a new bond to the proton of the Lewis acid, generating a carbocation. The H—Cl bond must break, giving its two electrons to Cl, forming Cl ¯. Because two electron pairs are involved, two curved arrows are needed. Acids and Bases