Acids and Bases

1. Acids and Bases Dr. Mai Yin Tsoi Georgia Gwinnett College

2. Definitions Water? Conjugate pairs pH Strong Acids Equilibrium Strikes Again!  Ka Ka and pH calculations Polyprotic Acids Bases Ions’ effect on pH Factors Affecting Acid Strength Lewis Acids and Bases Table of Contents

3. Qualitative Description • Acids – lemon juice, vinegar, acid rain • Sour • Litmus paper – red • Give off H+ ions in solution (Arrhenius) • Can “burn” you (react with proteins) • pH less than 7.0 • Bases – Windex, Drano, lye, • Bitter • Litmus paper – blue • Give off H+ ions in solution (Arrhenius) • Can “burn” you (react with proteins • pH greater than 7.0

4. Definitions (Bronsted-Lowry) • Acid – “generosity” • Can donate a “proton”, “H+” – call this “acidic” • Ex: HCl, HNO2, H2S, PH4+, HCO3- • Base – “selfishness” • Can receive a proton or H+ • NOT take—need non-bonded electrons • Ex: NaOH, NH3, ClO4-, CO, HPO42-

5. What about H2O? • Can act as an acid (generous): • H2O + NH3 NH4+ + OH- • Can act as a base (selfish): • H2O + HNO3  H3O+ + NO3- We call this “amphoteric” – can act as an either or base. Depends on what substance it’s mixed with! Ex: HCO3−, HSO4−

6. Acid in Water…What Happens? • Water acts as a Brønsted–Lowry base (selfish) and abstracts a proton (H+) from the acid (donated). • As a result, the conjugate base of the acid and a hydronium ion are formed.

7. Conjugate Acid-Base Pairs • Conjugate: “joined together as a pair” • Acid-base reaction: proton transfer

8. Acids HClO4 H3P AsH4+ HCO3- Bases HSO3- PO43- CO H2O Conjugate Bases HClO3- H2P- AsH3 CO32- Conjugate Acids H2SO3 HPO42- HCO+ H3O+ (hydronium ion) You Try!

9. How Strengths of the Pairs Relate • Strong acids- completely dissociate in water. So conjugate bases: • wussy!! • Weak acids-dissociate partially in water. So conjugate bases: • A bit stronger… • Animation on Acid Ionization • Possessive Boy/Girl Demo

10. pH Balanced for a Man/Woman… pH = −log [H3O+] • If pH = 3, what is [H+]? • If pH = 7, what is [H+]? • If pH = 12, what is [H+]? • If pH = 5.8, what is [H+]? • If [H+] = 3.5 x 104, what is pH?

11. pH and Molarity • Therefore, in pure water, pH = −log (1.0  10−7) = 7.00 • An acid has a higher [H3O+] than pure water, so its pH is <7 • A base has a lower [H3O+] than pure water, so its pH is >7.

12. .

13. How Do We Measure pH?

14. For more accurate measurements, one uses a pH meter, which measures the voltage in the solution. How Do We Measure pH Accurately?

15. Strong Acids • 7 strong acids: • HCl, HBr, HI, HNO3, H2SO4, HClO3, and HClO4 • Strong electrolytes - exist totally as ions in aqueous solution. • For monoprotic strong acids, [H3O+] = [acid]

16. Kc = [H3O+] [A−] [HA] HA(aq) + H2O(l) A−(aq) + H3O+(aq) Dissociation Constants • the equilibrium expression would be • This equilibrium constant is called the acid-dissociation constant, Ka.

17. [H3O+] [A−] [HA] Ka = Ka = [H3O+] [A−] [HA] Big or Small Ka? • Big Ka: • Strong acid • Mostly ions • Small Ka: • Weak acid • Mostly molecule

18. Got Ka? • The pH of a 0.10 M solution of formic acid, HCOOH, at 25°C is 2.38. Calculate Ka for formic acid at this temperature. • Remember def’n of pH • Remember def’n of Ka • What do they BOTH have in common? [H+]!

19. Calculating Ka from the pH • The pH of a 0.10 M solution of formic acid, HCOOH, at 25°C is 2.38. Calculate Ka for formic acid at this temperature. • To calculate Ka, we need the equilibrium concentrations of all three things. • We can find [H3O+], which is the same as [HCOO−], from the pH.

20. Calculating Ka from the pH pH = −log [H3O+] 2.38 = −log [H3O+] −2.38 = log [H3O+] 10−2.38 = 10log [H3O+] = [H3O+] 4.2  10−3 = [H3O+] = [HCOO−]

21. Calculating Ka from pH Now we can set up a table…

22. [4.2  10−3] [4.2  10−3] [0.10] Ka = Calculating Ka from pH = 1.8  10−4

23. Calculating pH from Ka Calculate the pH of a 0.30 M solution of acetic acid, HC2H3O2, at 25°C. HC2H3O2(aq) + H2O(l) H3O+(aq) + C2H3O2−(aq) Ka for acetic acid at 25°C is 1.8  10−5.

24. [H3O+] [C2H3O2−] [HC2H3O2] Ka = Calculating pH from Ka The equilibrium constant expression is

25. Calculating pH from Ka We next set up a table… We are assuming that x will be very small compared to 0.30 and can, therefore, be ignored.

26. 1.8  10−5 = (x)2 (0.30) Calculating pH from Ka Now, (1.8  10−5) (0.30) = x2 5.4  10−6 = x2 2.3  10−3 = x

27. Calculating pH from Ka pH = −log [H3O+] pH = −log (2.3  10−3) pH = 2.64

28. Polyprotic Acids • Have more than one acidic proton. • If the difference between the Ka for the first dissociation and subsequent Ka values is 103 or more, the pH generally depends only on the first dissociation.

29. Weak Bases Bases react with water to produce hydroxide ion.

30. [HB] [OH−] [B−] Kb = Weak Bases The equilibrium constant expression for this reaction is where Kb is the base-dissociation constant.

31. Weak Bases Kb can be used to find [OH−] and, through it, pH.

32. NH3(aq) + H2O(l) NH4+(aq) + OH−(aq) [NH4+] [OH−] [NH3] = 1.8  10−5 Kb = pH of Basic Solutions What is the pH of a 0.15 M solution of NH3?

33. pH of Basic Solutions Tabulate the data.

34. 1.8  10−5 = (x)2 (0.15) pH of Basic Solutions (1.8  10−5) (0.15) = x2 2.7  10−6 = x2 1.6  10−3 = x2

35. pH of Basic Solutions Therefore, [OH−] = 1.6  10−3M pOH = −log (1.6  10−3) pOH = 2.80 pH = 14.00 − 2.80 pH = 11.20

36. Kaand Kb Kaand Kb are related in this way: Ka Kb = Kw Therefore, if you know one of them, you can calculate the other.

37. X−(aq) + H2O(l) HX(aq) + OH−(aq) Reactions of Anions with Water • Anions are bases. • As such, they can react with water in a hydrolysis reaction to form OH− and the conjugate acid:

38. Reactions of Cations with Water • Cations with acidic protons (like NH4+) will lower the pH of a solution. • Most metal cations that are hydrated in solution also lower the pH of the solution.

39. Reactions of Cations with Water • Attraction between nonbonding electrons on oxygen and the metal causes a shift of the electron density in water. • This makes the O-H bond more polar and the water more acidic. • Greater charge and smaller size make a cation more acidic.

40. Effect of Cations and Anions • An anion that is the conjugate base of a strong acid will not affect the pH. • An anion that is the conjugate base of a weak acid will increase the pH. • A cation that is the conjugate acid of a weak base will decrease the pH.

41. Effect of Cations and Anions • Cations of the strong Arrhenius bases will not affect the pH. • Other metal ions will cause a decrease in pH. • When a solution contains both the conjugate base of a weak acid and the conjugate acid of a weak base, the affect on pH depends on the Kaand Kbvalues.

42. Factors Affecting Acid Strength • The more polar the H-X bond and/or the weaker the H-X bond, the more acidic the compound. • Acidity increases from left to right across a row and from top to bottom down a group.

43. Factors Affecting Acid Strength In oxyacids, in which an OH is bonded to another atom, Y, the more electronegative Y is, the more acidic the acid.

44. Factors Affecting Acid Strength For a series of oxyacids, acidity increases with the number of oxygens.

45. Factors Affecting Acid Strength Resonance in the conjugate bases of carboxylic acids stabilizes the base and makes the conjugate acid more acidic.

46. Lewis Acids • Lewis acids are defined as electron-pair acceptors. • Atoms with an empty valence orbital can be Lewis acids.

47. Lewis Bases • Lewis bases are defined as electron-pair donors. • Anything that could be a Brønsted–Lowry base is a Lewis base. • Lewis bases can interact with things other than protons, however.