1. Acids and Bases
Chapter 15

2. I. Properties of Acids & Bases

3. Example Acids
Citric acid
Stomach acid (HCl)
Vinegar
Ascorbic acid

4. Properties of Acids Taste sour React with metals Often corrosive
Turn blue litmus paper red
Watch someone eat a lemon
Reacts with metals to form hydrogen gas

5. Common Acids

6. Example Bases
Ammonia
Baking soda
Antacids
Soaps
Cleaners

7. Properties of Bases Taste bitter Feel slippery Can burn
Turn red litmus paper blue
Taste of vinegar, slippery like soap

8. Common Bases

9. II. Acid- Base Theories

10. Arrhenius Theory of Acids
A chemical compound that increases the concentration of hydrogen ions (H+) in aqueous solution
Hydronium ion (H3O+) – formed from an aqueous acid solution
Example: HCl + H2O  H3O+ + Cl-

11. Arrhenius Theory of Bases
A substance that increases the concentration of hydroxide ions (OH-) in aqueous solution
Example: NaOH(aq)  Na+(aq) + OH–(aq)

12. Brønsted–Lowry Acid–Base Theory
Acid: proton (H+) donor
HCl + NH3  NH4+ + Cl–
HCl donates a H+ to NH3
H+ does not exist by itself

13. Bases: accept a proton
H2O + NH3  NH4+ + OH–
NH3 accepts a H+ from H2O

14. Acid/Base Conjugates
Brønsted–Lowry theory defines acids and bases in terms of proton (H+) transfer.
A Brønsted–Lowry acid is a proton donor.
A Brønsted–Lowry base is a proton acceptor.
The conjugate base of an acid is the acid minus the proton it has donated.
The conjugate acid of a base is the base plus the accepted proton.

15. Practice problems
Identify the acid, base, conjugate acid, conjugate base, and conjugate acid-base pairs:
HC2H3O2(aq) + H2O(l)  C2H3O2–(aq) + H3O+(aq)
acid
base
conjugate base
conjugate acid
conjugate acid-base pairs
OH –(aq) + HCO3–(aq)  CO32–(aq) + H2O(l)
base
acid
conjugate base
conjugate acid
conjugate acid-base pairs

16. III. Acids & Bases with Water

17. Amphiprotic A substance that can either donate or accept a proton
H2O most common amphiprotic substance
Reacts with base: H2O + NH3  NH4+ + OH–
Reacts with acid: H2O + HCl  Cl– + H3O+

18. Aqueous Acids and Bases
Water as both an acid and a base
Arrhenius H2O  H+ + OH-
Brønsted-Lowry H2O + H2O  H3O+ + OH-

19. Strong and Weak Acids Strong acid: acid completely reacts with water
Completely ionizes in water
Weak acid: acid reacts only slightly with water
Partially ionizes in water
Common strong acids: HCl, HNO3, H2SO4
Most other acids are weak acids
Strong acids are strong electrolytes

20. Strong and Weak Bases Strong base: base completely ionizes in water
Weak base: only slightly ionizes in water
Common strong bases: NaOH and KOH
True for most Group 1A and 2A hydroxides
Common weak base: NH3

21. Aqueous Acids and Bases
Acid in water
Ionic compounds
Strength based on concentration of H+ in solution
Base in water
Strength based on concentration of OH- in solution
Ions break apart in water

22. Aqueous Acids and Bases
Pure water is perfectly neutral
Not the slightest bit acidic or basic
Water ionization is reversible H2O + H2O  H3O+ + OH-
There exists a dynamic equilibrium
Simultaneous ionization and recombination
Rate of the forward reaction = rate of the reverse reaction

23. Neutralization Reaction of an acid and a base to form water and a salt
HCl + NaOH  NaCl + H2O

24. IV. pH Scale

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

26. pH in Common Solutions

27. Measurement of Acidity
H2O + H2O  H3O+ + OH-
The self-ionization of water at 25˚C
Equal amounts of H+ and OH- are produced
Equal concentrations of 1 x 10-7 moles/liter
Neutral water: [H+ ] = M or 1 x 10-7M
[OH-] = M or 1 x 10-7M
Ionization constant of water – Kw
Kw = [H3O+][OH-] = (1 x 10-7M)(1 x 10-7M) = 1 x 10-14M
pH represents the power of the Hydrogen
Negative logarithm of the H+ concentration
pH = -log [H+ ]
= -log 10-7
= -(-7)
pH = +7

28. What is the pH of 1 x 10-3M of HCl?

29. pH = 3

30. What is the pH for 1 x 10-4M of NaOH

31. pH = 10

32. What is the pH of a solution if the [H3O+] is 3.4 x 10-5M?

33. pH = -log [H3O+] = -log (3.4 x 10-5) = 4.47

34. Calculating the concentration from the pH [H3O+] = antilog (-pH) What is the hydronium concentration of an acid solution having a pH of 3?

35. What is the hydronium concentration of an acid solution having a pH of 5.3?

36. Indicators Compounds whose colors are sensitive to pH
Methyl orange – changes colors with a pH between 3.1 and 4.4
Phenolphthalein – changes colors with a pH between 8.0 and 10.0

37. Nature’s Indicator
Many flowers and plants contain acid/base indicators which change color in response to a change in pH
Chemical structure changes when pH changes
Soil pH and anthocyanins

38. Titration
The controlled addition and measurement of the amount of a solution of known concentration required to react completely with a measured amount of a solution of unknown concetration

39. Antacids and Acid Indigestion
Antacids are used to decrease the production of excess stomach acid
Active ingredients
Hydroxides or carbonates
React with stomach acid (HCl)
Too much of a good thing!
Acid rebound

40. If a solution ionizes almost completely to form OH- ions, what is the substance?

41. What is a conjugate acid/base pair in the following equation
What is a conjugate acid/base pair in the following equation? HCN + H2O  H3O + CN-

42. When the hydroxide ion concentration is greater than the hydronium ion concentration, is the solution an acid or a base?

43. What is the pH of a 3.2 x 10-6M?

44. What is the pH of a 1.0 x 10-3M solution of NaOH?