1. Organic Chemistry PRINCIPLES AND MECHANISMS Chapter 6: Lecture PowerPoint The Proton Transfer Reaction

2. A proton transfer reaction (or a Brønsted–Lowry acid–base reaction) is one in which a Brønsted–Lowry base reacts with a Brønsted–Lowry acid. Brønsted–Lowry bases accept protons (H + ), and Brønsted– Lowry acids donate the protons. A proton transfer reaction consists of a single elementary step, in which the bonds break and form simultaneously. 6.1 An Introduction to Reaction Mechanisms: The Proton Transfer Reaction and Curved Arrow Notation

3. Curved Arrow Notation This notation (also called arrow pushing) keeps account of the valence electrons involved in the reaction mechanism.

4. Curved Arrow Notation continued… Bond breaking and bond formation involve only valence electrons. Curved arrows show the appropriate movement of those electrons.

5. 6.2 Chemical Equilibrium and the Equilibrium Constant, K eq A proton transfer reaction between HCl and HO⁻ will take place readily (i.e., strong acid/strong base reaction). Essentially no proton transfer will occur between HO⁻ and NH 3 (strong base/weak acid reaction). A reaction’s tendency to form products is described by its equilibrium constant, K eq.

6. Proton Transfer and K eq For a proton transfer reaction between HA (the acid) and B⁻ (the base), the equilibrium constant expression is written as:

7. Acid Strengths: K a and pK a The strength of an acid can be thought of as its ability to donate a proton. Acid strengths can be obtained experimentally. The acidity constant, K a, eliminates H 2 O from the equation. When comparing two acids, the one with the larger K a value is the stronger acid.

8. Acid Strengths: K a and pK a continued… Chemist often work with pK a instead of K a due to K a ’s immensely large range of values. pK a = -log(K a ) As the value of K a increases, pK a decreases. Compounds with low pK a values are more acidic than compounds with high pK a values. Each difference of 1 in pK a values represents a factor of 10 difference in acid strength.

9. Predicting the Outcome of a Proton Transfer Reaction Using pK a Values Proton transfer reactions favor the side opposite the stronger acid.

10. The Leveling Effect Even if a solvent does not act as an acid or a base in an intended proton transfer reaction, it can limit the existence of certain acids and bases in solution. These restrictions are the result of what is called the leveling effect.

11. The Leveling Effect and Other Solvents The strongest acid that can exist in solution to any appreciable concentration is the protonated solvent. The strongest base that can exist in solution is the deprotonated solvent. In water, no acid stronger than H 3 O + and no base stronger than HO⁻ can exist to any appreciable extent.

12. Le Châtelier’s Principle and pH Le Châtelier’s principle states that if a reaction at equilibrium experiences a change in reaction conditions (e.g., concentrations, temperature, pressure, or volume), then the equilibrium will shift to counteract that change.

13. In water, increasing the concentration of H 3 O + will cause the equilibrium to shift toward reactants. At the new equilibrium, there will be more HA and less A⁻, so the percent dissociation of HA will decrease. Additionally, decreasing the concentration of H 3 O + will cause the equilibrium to shift toward products, which results in an increase in the percent dissociation of HA. The concentration of H 3 O + in solution is reported as pH = - log[H 3 O + ]. Le Châtelier’s Principle and pH continued…

14. Henderson–Hasselbalch Equation The Henderson-Hasselbalch equation is a common general chemistry equation used for dealing with buffers. At 50% dissociation of the acid, the equilibrium concentrations of A⁻ (aq) and HA(aq) are equal. This makes the second term in Equation 6-16 equal to zero, in which case pH = pK a.

15. Percent Dissociation of an Acid with pK a = 5 as a Function of pH

16. 6.3 Thermodynamics and Gibbs Free Energy Equilibrium constant is also related to the standard Gibbs free energy difference (  G o rxn ) between reactants and products. The naught ( o ) signifies standard conditions (298 K, 1 atm, and the concentrations of all solutions are 1 mol/L)

17. 6.3 Thermodynamics and Gibbs Free Energy continued… Because the gas constant (R) and temperature (T) are both positive, a negative value for  G o rxn results when K eq > 1. Conversely, a positive value for  G o rxn means that K eq < 1. When K eq 1, products are favored.

18. Enthalpy and Entropy The  G o rxn can be expressed in terms of the reaction’s enthalpy change (  H o rxn ) and entropy change (  S o rxn ). The  H o rxn term is the standard enthalpy difference between the reactants and products. At constant pressure,  H o rxn equals the heat absorbed or released by the reaction. If  H o rxn is “+”, then the reaction absorbs heat (endothermic). If  H o rxn is “-”, then the reaction releases heat (exothermic).

19. Enthalpy and Entropy continued…  S o rxn is the standard entropy difference between the reactants and products.  S o rxn is often thought of as a “measure of disorder.”  S o rxn for a proton transfer reaction is usually very small.

20. The Reaction Free Energy Diagram In a reaction free-energy diagram, Gibbs free energy is plotted as a function of the reaction coordinate. A reaction coordinate is a variable that corresponds to the changes in geometry, on a molecular level, as reactants are converted into products.

21. 6.4 Strategies for Success: Functional Groups and Acidity One way to estimate the unknown pK a value of a compound is to observe any structural similarities to a compound that is already tabulated. The acidity (pK a ) of a compound is mostly governed by the functional group on which the acidic proton is found. Acidity is relatively independent of the molecule’s carbon skeleton, as observed by the structures at the right.

22. Factors That Affect pK a Nearby structural features, such as a highly electronegative substituent or adjacent double bond, can alter the acidity significantly. Trichloroacetic acid (Cl 3 CCO 2 H; pK a = 0.77) is a stronger acid than acetic acid (H 3 CCO 2 H; pK a = 4.75). Phenol is 10 6 times stronger than ethanol!

23. 6.5 Relative Strengths of Charged and Uncharged Acids: The Reactivity of Charged Species There are trends that dictate the strength of an acid. A proton is significantly more acidic when it is attached to a positively charged atom than when that atom is uncharged. The pK a of H 3 O + is -1.7, whereas that of H 2 O is 15.7. The pK a of H 4 N + is 9.4, whereas that of H 3 N is 36.

24. 6.5 Relative Strengths of Charged and Uncharged Acids: The Reactivity of Charged Species continued… The stronger acidity of positively charged acids is a reflection of lower stability of charged species than similar uncharged species.

25. The Reactivity of Charged Species H 3 O + and HO⁻ products, which are charged, are less stable than the uncharged reactants by 80 kJ/mol.

26. 6.6 Relative Acidities of Protons on Atoms with Like Charges The farther to the right an atom appears in the periodic table, the more acidic the protons that are attached to it. pK a : CH 4 > NH 3 > H 2 O > HF Negative charge generated by loss of H + is better stabilized on the more electronegative atom.

27. 6.6 Relative Acidities of Protons on Atoms with Like Charges continued…

28. Protons on Different Atoms in the Same Column of the Periodic Table The farther down a column an atom appears in the periodic table, the more acidic the protons that are attached to it. pK a : HF > HCl The stability of an anion increases when the negative charge is on an atom farther down a column of the periodic table. The negative charge is more distributed over the larger atoms, which are found as you go down the column (i.e., principal quantum number increases).

29. Hybridization of the Atom to Which the Proton Is Attached Ethane, ethene, and ethyne are all hydrocarbons, yet differ in hybridization of their carbon atoms (sp 3, sp 2, and sp, respectively). These three hydrocarbons are also uncharged acids. Their differences in acidity must be caused by differences in the stability of their conjugate bases. In other words, HCΞC⁻ is more stable than H 2 C=CH⁻, which is more stable than H 3 C-CH 2 ⁻. The C atom bearing the negative charge in HCΞC⁻ is sp hybridized (highest effective electronegativity = greatest stability).

30. Effects from Adjacent Double and Triple Bonds: Resonance Effects Ethanoic acid and ethanol are both uncharged acids yet ethanoic acid is a considerably stronger acid. Ethanoic acid has a C=O bond adjacent to the acidic OH group, and is more acidic than ethanol (CH 3 CH 2 OH) by about 11 pK a units. One of the main reasons ethanoic acid is so much more acidic than ethanol has to do with resonance effects.

31. Effects from Adjacent Double and Triple Bonds: Resonance Effects continued… Ethanoic acid’s conjugate base is highly stabilized by resonance, whereas ethanol’s negative charge is localized on the O atom.

32. Resonance Stabilization The more delocalized a negative charge, the more stable the conjugate base becomes. Resonance stabilization generally increases as the number of atoms over which a charge is delocalized increases.

33. Resonance Stabilization: H 2 SO 4 vs Ethanoic Acid Sulfuric acid offers a greater number of electronegative atoms to delocalize its negative charge (3 total) as compared to ethanoic acid (2 total). This additional delocalized site helps explain why sulfuric acid has a pK a = -9, while ethanoic acid’s pK a = 4.75!

34. Effects from Nearby Atoms: Inductive Effects The presence of nearby electronegative atoms increases the stability of the conjugate base. 2-Chloroethanol (ClCH 2 CH 2 OH; pK a = 12.9) is more acidic than ethanol (CH 3 CH 2 OH; pK a = 16). The O atom bears less of a negative charge in ClCH 2 CH 2 O⁻ than it does in CH 3 CH 2 O⁻.

35. Inductive Effects Cl is electron withdrawing in comparison to H. The electron density on the left-most C atom of 2- chloroethanol is shifted toward the Cl atom. To compensate for this shift, it would draw electron density away from other atoms in which it is bonded. This effect is repeated down the chain until the electron density is pulled from the O atom. This is called induction.

36. Inductive Effects continued…

38. Anion and Cation Stabilization Most atoms are more electronegative than H so their presence is usually electron withdrawing. Some atoms are actually electron donating, when compared to H (example: the Si atom). Alkyl groups are also electron donating.

39. Anion and Cation Stabilization continued… Electron-donating groups stabilize nearby cations. The inductive effects of alkyl groups stabilizes the nearby positive charge. Carbocations (C+) are increasingly stabilized by the addition of more alkyl groups. The alkyl groups help diminish the positive charge by inductive effects.

40. Carbocations and Their Stability Carbocations are distinguished by their degree of alkyl substitution. Carbocation stability increases in the following order: Methyl < 1 o < 2 o < 3 o

41. 6.7 Strategies for Success: Ranking Acid and Base Strengths—The Relative Importance of Effects on Charge To predict the relative stabilities of two species, the following questions must be asked in the following order. 1.Do the species have different charges? A charged species is usually more reactive than one that bears no formal charge. 2.Do the charges appear on different atoms? The size, electronegativity, and hybridization of the atom on which the charge is located affects stability.

42. 6.7 Strategies for Success: Ranking Acid and Base Strengths—The Relative Importance of Effects on Charge continued… 3.Are the charges delocalized differently through resonance? All else being equal, stability increases as the number of atoms over which a charge is shared increases. 4.Are there differences in inductive effects? Electron-withdrawing groups stabilize nearby negative charges but destabilize nearby positive charges. Electron-donating groups stabilize nearby positive charges but destabilize nearby negative charges.

43. Acid Strength Because pKa is defined in terms of each acid’s reaction with water, we begin by writing out their reactions.

44. Free Energy Diagrams for the Proton Transfer Reactions The acid in Equation 6-24c (red) is charged, whereas those in Equation 6-24a, 6- 24b, and 6-24d are uncharged. In Figure 6-13, therefore, the reactants in Equation 6- 24a, 6-24b, and 6-24d are lower in energy than the reactants in Equation 6-24c. Because the acids in Equation 6-24a, 6-24b, and 6-24d are all uncharged, the remaining three tie- breaking questions cannot be used to distinguish their relative stabilities; their stabilities are assumed to be similar.

45. Free Energy Diagrams for the Proton Transfer Reactions continued…

46. Moving to the product side, notice first that H 3 O + can be ignored in each case, so the tie-breaking questions should be applied just to the conjugate bases. Acid strength therefore increases in the order: A < B < C < D Free Energy Diagrams for the Proton Transfer Reactions continued…

47. Conjugate Base Strengths If you know the order of acid strengths, then it is straightforward to determine the order of base strengths.

48. 6.8 Strategies for Success: Determining Relative Stabilities of Resonance Structures Although the resonance hybrid is an average of all the resonance contributors, not all resonance contributors are weighted equally. The resonance hybrid looks most like the lowest energy (most stable) resonance structure. The second resonance structure on the right is less stable for having formal charges. Therefore, that structure makes a smaller contribution to the resonance hybrid. Case #1 Charged verses Uncharged Resonance Form

49. 6.8 Strategies for Success: Determining Relative Stabilities of Resonance Structures continued… The main difference in the following example is where the formal charge is localized. Negative charge on the more electronegative atom is more stable. Therefore, the resonance structure on the left having the “ ˗ ” charge on the carbon atom (more electropositive atom), is less stable. Resonance structure on the left contributes less to the resonance hybrid. Case #2 Negative charge on different atoms.

50. 6.8 Strategies for Success: Determining Relative Stabilities of Resonance Structures continued… The resonance structure which possesses the “+” charge on the more substituted C atom is the more stable one. Again, the greater the degree of alkyl substitution, the more stable the carbocation. Case #3 Positive charge on different carbon atoms.

51. 6.10 The Structure of Amino Acids in Solution as a Function of pH Amino acids have both the amino and the carboxyl group. This form never dominates in aqueous solution. The carboxyl group is weakly acidic and the amine group is weakly basic. The form that the amino acid takes depends upon the pH of the solution.

52. The Zwitterion Under highly acidic (pH < 2) conditions, the weakly basic N atom is protonated. The resulting species, which bears an overall charge of +1, is shown on the left in the following equation. When the pH of the solution has risen significantly above 2, the OH proton is lost and the dominant form is the middle species (this makes a zwitterion). When the pH of the solution is significantly above 9 or 10 (the pK a of the protonated amine), the anion on the right dominates.

53. 6.11 Electrophoresis and Isoelectric Focusing Electrophoresis, one of the most common ways of separating a mixture of amino acids or proteins. A mixture is spotted on a gel or strip of paper that has been buffered to a specific pH. A high voltage (50–1,000 V) is applied by the anode and cathode ends.

54. The Isoelectric Point (pI) Any ion that has a net positive charge at that pH will migrate toward the negatively charged cathode. Any ion having a net negative charge will migrate toward the positively charged anode. If the net charge is zero, the species will not move. An amino acid’s isoelectric pH, or isoelectric point (pI), is the pH at which the substance has a charge of zero. The pI is unique for each of the 20 amino acids (pI = 6.07 for glycine; pI = 2.98 for aspartic acid; pI = 9.74 for asparagine).

55. Summary and Conclusions Proton transfer reactions occurs when a Brønsted acid donates an H + to a Brønsted base. Curved arrow notation shows how the valence electrons are involved in the proton transfer reaction. The equilibrium constant reflects the tendency to form products. The acidity constant, K a, reflects an acid’s tendency to donate a proton. Electronegativity, charge, hybridization, and resonance all affect a compound’s acid strength.

56. Summary and Conclusions continued… A solvent’s leveling effect dictates the maximum strength of an acid or a base that can exist in solution. The percent dissociation of an acid increases as the pH of the solution increases. Functional groups mostly govern a compound’s pKa. If  G o rxn is negative, the reaction tends to be spontaneous. If  G o rxn is positive, it tends to be nonspontaneous.

57. Summary and Conclusions continued… For two ions in which the formal charge is on an atom of the same element, hybridization governs stability. Resonance effects can stabilize a charged species. Inductive effects can affect the stability of a charged species by shifting electron density through covalent bonds.