1. Bettelheim, Brown, Campbell and Farrell Chapter 9
Acids and Bases
Bettelheim, Brown, Campbell and Farrell
Chapter 9

2. Arrhenius Acids and Bases
acid: a substance that produces H3O+ ions aqueous solution
When HCl dissolves in water, its reacts with water to give hydronium ion and chloride ion

3. Arrhenius Acids and Bases
base: a substance that produces OH- ions in aqueous solution
other bases are not hydroxides; these bases produce OH- by reacting with water molecules

4. Solutions Acidic Solution: H3O+ > OH- (low pH) Basic (Alkaline)
(high pH)
Neutral Solution: H3O+ = OH-
(pH ~ 7)

6. Acid and Base Strength
Strong acid: one that reacts completely or almost completely with water to form H3O+ ions
Strong base: one that reacts completely or almost completely with water to form OH- ions

7. Strong Acids & Bases
All others are weak acids or bases

8. Acid and Base Strength
Weak acid: a substance that dissociates only partially in water to produce H3O+ ions
Weak base: a substance that dissociates only partially in water to produce OH- ions

9. Brønsted-Lowry Acids & Bases
Acid: a proton donor
Base: a proton acceptor
Acid-base reaction: a proton transfer reaction
Conjugate acid-base pair: any pair of molecules or ions that can be interconverted by transfer of a proton

10. Conjugate acid base pair differ only by a proton H+

11. Brønsted-Lowry Acids & Bases
Brønsted-Lowry definitions do not require water as a reactant

12. Conjugate Acids and Bases
HA B ↔ BH A-
acid base conjugate conjugate
acid of B base of HA
An acid will react to form its conjugate base.
A base will react to form its conjugate acid.
Conjugate acid-base pairs differ only by a hydrogen

13. Conjugate Acids and Bases
HA B ↔ BH A-
acid base conjugate conjugate
acid of B base of HA
An acid will react to form its conjugate base.
A base will react to form its conjugate acid.
Conjugate acid-base pairs differ only by a hydrogen

14. Conjugate Acid-Base Pairs
If acid is very strong, its corresponding conjugate base is very weak.
Stronger acid ionizes more completely, so the base will not attract the H3O+ well
Example:
HCl Cl-

15. Conjugate Acid-Base Pairs
If base is stronger, its corresponding conjugate acid is weaker.
Stronger base accepts H+ more easily, so the acid will not donate H+ well
Examples: OH HOH
PO33- HPO42-

16. Fig. 9.2
Weak
Strong
Strong
Weak

18. Water as an Acid and a Base
Water is amphoteric:
HC2H3O2 + H2O H3O+ + C2H3O2-
acid base acid base
H2O + NH NH4+ + OH-
acid base acid base .. .. .. 14

19. Acid-Base Equilibria
For weak acids, significant amounts of both the acid and its conjugate base will be present and form an equilibrium
Strong Acid
Weak Acid

20. Equilibrium lies on the side of the weaker acid and weaker base.
Acid-Base Equilibria
What if the base is not water? How can we determine which are the major species present?
Equilibrium lies on the side of the
weaker acid and weaker base.

22. Example: Predict the position of equilibrium in this
Acid-Base Equilibria
The position of this equilibrium lies to the right—
Formation of weaker acid and weaker base is favored
Example: Predict the position of equilibrium in this
acid-base reaction

23. Bases: OH- stronger than HCO32- Acids: H2CO3 stronger than H2O
Weaker acid
Weaker base
Can see from table:
Bases: OH- stronger than HCO32-
Acids: H2CO3 stronger than H2O
Right side favored—Equilibrium to the right

24. Acid Ionization Constants
The equilibrium constant, Keq, is
Treat [H2O] as a constant equal to 1000 g/L or 55.5 mol/L
combine the two constants to give a new constant, which we call an acid ionization constant, Ka

25. Acid Ionization Constants: Ka
p = -log
Acetic Acid: Ka for acetic acid is 1.8 x 10-5
pKa for acetic acid is 4.75

26. Acid Ionization Constants: Ka
p = -log
Weak acid has the small Ka, but large pKa

28. Self-Ionization of Water
[H2O] as a constant = 55.5 mol/L

29. Self-Ionization of Water
Ion product of water, Kw = 1.0 x 10-14
Product of [H3O+] and [OH-] in any aqueous solution is equal to 1.0 x 10-14

30. Self-Ionization of Water
Example: add mole of HCl to 1 liter of pure water, in this solution, [H3O+] is or 1.0 x What is hydroxide ion concentration?
pH = -log [H3O+] p = -log
pOH = -log [OH-]

31. pH Scale By definition: pH = - log [H3O+] p = “- log”
Brackets used to show concentration (M)
Scale ranges from 0 to 14

32. pH scale Range from 0 to 14 pH = -log [H3O+] pH < 7 acidic
pH = 7 neutral
pH > 7 basic

33. Solutions Acidic Solution: H3O+ > OH- Basic (Alkaline)
Neutral Solution: H3O+ = OH-

34. pH Scale Typical values range from 0 to 14 pH = 7 – neutral
pH > 7 – basic
pH < 7 – acidic

35. pH
pH of some common materials

36. pH of Salt Solutions NaCH3COO = Sodium Acetate In water:
NaCH3COO → Na CH3COO-
Na+ + OH- → NaOH (strong base)
CH3COO H+ → CH3COOH (weak acid)
Strong base + weak acid = basic salt

37. pH of Salt Solutions NH4Cl Ammonium chloride
NH OH- → NH4OH (weak base)
H Cl- → HCl (strong acid)
Strong acid + weak base = acidic salt

38. Neutralization Reactions
Also known as Acid-Base Reactions
H3O OH- → H2O
Acid Base → Water
Neutralization is Special Case of Double Replacement Reaction--- Water is a product
HCl + NaOH → NaCl H2O
Acid Base → Salt Water

39. Neutralization Reaction
Acid and Base react with each other to form water and a salt
Frequently use a pH indicator to show when end point has been reached
End point is pH at which [H3O+] = [OH-]
and color chosen indicator changes

40. Titration
Titration:
Volume of a solution of known concentration is added to a solution of unknown concentration.
Measure amount of known solution needed to react exactly with original material - at endpoint

41. Equivalence point:
Amount of acid = Amount of base
End point:
Indicator color changes

42. Fig. 9.6

43. Acid-Base Titrations
Use M H2SO4 to determine the concentration of a NaOH solution
HOW?

44. Example:
Add NaOH (a base) to a solution of acid that contains the pH indicator phenolphthalein
Phenolphthalein is colorless in acidic solution and is pink in basic solution
When enough NaOH is added so that it neutralizes all of the acid, any additional NaOH makes the solution basic and it turns pink

45. What is the concentration of an acid solution if 12. 54 mL of 0
What is the concentration of an acid solution if mL of M NaOH neutralizes mL of acid?
At endpoint: mol Acid = mol Base
MaVa = MbVb
Ma = MbVb = ( M) ( mL)
Va mL
= M

46. pH indicator (acid-base indicator): substance that turns color when the H3O+ (acid) concentration changes

48. pH indicators may be extracted from many natural products
lichens, apple skins, blueberries,
red cabbage, etc.
pH indicators may be in many forms: embedded in paper (pH or litmus paper)
liquid

49. Buffer Solutions Contain a weak acid and its conjugate base
Contain a weak base and its conjugate acid
Two species differ only by an H+
Add extra acid—the conjugate base will react to remove the added H+
Add extra base—the conjugate acid will react to remove the added OH-
Extra acid or base is removed, so pH will remain relatively constant

50. pH Buffers
add a strong acid, such as HCl, adds H3O+ ions react with acetate ions and are removed from solution
add a strong base, such as NaOH, adds OH- ions react with acetic acid and are removed from solution

51. pH Buffers
Consider a phosphate buffer of 0.10 mole of NaH2PO4 (a weak acid) and 0.10 mole of Na2HPO4 (the salt of its conjugate base)

52. Illustration of Buffer Effects
What is the pH of a solution 0.1 M in formic acid (HCOOH) and 0.10 M in HCOONa when Ka=1.8 x 10-4?
What is the pH of the solution after 0.03 mol of NaOH is added to 1.0 L of the buffer?
[H3O+]= Ka [HA]/[A-] [H3O+] = 1.8 x 10-4 (0.10)/(0.10)
pH= 3.8
[H3O+] = 1.8 x 10-4 (0.07)/(0.10)
pH = 3.9

53. Blood Buffers The average pH of human blood is 7.4
The body uses three buffer systems
carbonate buffer: H2CO3 and its conjugate base, HCO3-
phosphate buffer: H2PO4- and its conjugate base, HPO42-
proteins: discussed in Chapter 21

54. Henderson-Hasselbalch Eqn
Henderson-Hasselbalch equation: shows mathematical relationship between
pH,
pKa of the weak acid, HA
concentrations HA, and its conjugate base, A-