1. ACIDS and BASES Unit 10, Chapter 19

2. pH indicators
pH indicators are valuable tool for determining if a substance is an acid or a base.
The indicator will change colors in solution.

3. Things to use…
pH meter will indicate the numeric value of acid or base based on the pH range
Chemical indicators: phenolphthalein, universal indicator…
Natural indicators: poinsettia, red cabbage juice…

4. DEMO: Red Cabbage Juice
Name of Product
Prediction: Acid or Base?
Indicator Color Change
Acid or Base?
HCl
Sprite
Aspirin
Antacid
Vinegar
Window Cleaner
Baking soda
NaOH

5. Properties of Acids and Bases
Have a sour taste
Change the color of many indicators
Are corrosive (react with metals)
Neutralize bases
Conduct an electric current
BASES
Have a bitter taste
Change the color of many indicators
Have a slippery feeling
Neutralize acids
Conduct an electric current

6. The Arrhenius Theory of Acids and Bases

7. Arrhenius Theory of Acids and Bases:
an acid contains hydrogen and ionizes in solutions to produce H+ ions:
HCl  H+(aq) + Cl-(aq)

8. Arrhenius Theory of Acids and Bases:
a base contains an OH- group and ionizes in solutions to produce OH- ions:
NaOH  Na+(aq) + OH-(aq)

9. Neutralization
Neutralization: the combination of H+ with OH- to form water.
H+(aq) + OH-(aq)  H2O (l)
Hydrogen ions (H+) in solution form hydronium ions (H3O+)

10. In Reality… H+ + H2O  H3O+ Hydronium Ion
(Can be used interchangeably with H+)

11. Commentary on Arrhenius Theory…
One problem with the Arrhenius theory is that it’s not comprehensive enough. Some compounds act like acids and bases that don’t fit the standard definition.

12. Bronsted-Lowry Theory of Acids & Bases

13. Bronsted-Lowry Theory of Acids & Bases:
An acid is a proton (H+) donor
A base is a proton (H+) acceptor

14. for example… Proton transfer HCl(aq) + H2O(l)  H3O+(aq) + Cl-(aq)
Base
Acid

15. another example… Water is a proton donor, and thus an acid.
CONJUGATE BASE
ACID
NH3(aq) + H2O(l)  NH4+ (aq) + OH- (aq)
BASE
CONJUGATE ACID
Ammonia is a proton acceptor, and thus a base

16. Amphoteric Substances
A substance that can act as both an acid and a base (depending on what it is reacting with) is termed amphoteric.
Water is a prime example.

17. Conjugate acid-base pairs
Conjugate acid-base pairs differ by one proton (H+)
A conjugate acid is the particle formed when a base gains a proton.
A conjugate base is the particle that remains when an acid gives off a proton.

18. Examples: In the following reactions, label the conjugate acid-base pairs:
H3PO4 + NO2-  HNO2 + H2PO4-
CN- + HCO3-  HCN + CO32-
HCN + SO32-  HSO CN-
H2O + HF  F- + H3O+
acid
base
c. acid
c. base
base
acid
c. acid
c. base
acid
base
c. acid
c. base
base
acid
c. base
c. acid

19. SUMMARY OF ACID-BASE THEORIES
Theory
Acid Definition
Base Definition
Arrhenius Theory
Any substance which releases H+ ions in water solution.
Any substance which releases OH- ions in water solution
Brǿnsted-Lowry Theory
Any substance which donates a proton.
Any substance which accepts a proton.

20. Strength of Acids and Bases
A strong acid dissociates completely in sol’n:
HCl  H+(aq) + Cl-(aq)
A weak acid dissociates only partly in sol’n:
HNO2  H+(aq) + NO2-(aq)
A strong base dissociates completely in sol’n:
NaOH  Na+(aq) + OH-(aq)
A weak base dissociates only partly in sol’n:
NH3(aq) + H2O(l)  NH4+(aq) + OH-(aq)

21. Acid-Base Reactions H+(aq) + OH-(aq)  H2O(l)
Neutralization reactions: reactions between acids and metal hydroxide bases which produce a salt and water.
H+ ions and OH- ions combine to form water molecules:
H+(aq) + OH-(aq)  H2O(l)

22. Buffered Solutions
A solution of a weak acid and a common ion is called a buffered solution.

23. Thus, the solution maintains it’s pH in spite of added acid or base.
Pg. 620 fig 19.27

24. Demo: buffered solution

25. Demo: tap water vs. dH2O
Both waters have Universal indicator in them (= pH indicator (changes color in the presence of ions), which is a type of weak acids)
The water will change pH, and therefore COLOR (which helps us determine if a solution is acidic or basic) with the addition of HCl (acid) and NaOH (base)

26. Universal Indicator Color Chart PAGE 602 fig 19.13
pH scale
Acid Neutral Base

27. Why does it take more drops of acid or base to make the tap water change color than it does for the distilled water?
What is distilled water made of? What is tap water made of?

28. pH and pOH
Pg. 596 (in section 19.2)

29. Ionization of water
Experiments have shown that pure water ionizes very slightly:
2H2O  H3O+ + OH-
Measurements show that:
[H3O+] = [OH-]=1 x 10-7 M
Pure water contains equal concentrations of H3O+ + OH-, so it is neutral.

30. pH pH = -log [H3O+] pH = -log [H+]
pH is a measure of the concentration of hydronium ions in a solution.
pH = -log [H3O+] or
pH = -log [H+]

31. Example: What is the pH of a solution where [H3O+] = 1 x 10-7 M?
pH = -log [H3O+]
pH = -log(1 x 10-7)
pH = 7

32. Example: What is the pH of a solution where [H3O+] = 1 x 10-5 M?
pH = -log [H3O+]
pH = -log(1 x 10-5)
pH = 5
When acid is added to water, the [H3O+] increases, and the pH decreases.

33. Example: What is the pH of a solution where [H3O+] = 1 x 10-10 M?
pH = -log [H3O+]
pH = -log(1 x 10-10)
pH = 10
When base is added to water, the [H3O+] decreases, and the pH increases.

34. The pH Scale PAGE 598 Table 19.5 & fig 19.10
*You must use pH to determine if something is acidic, basic or netural (not pOH)
Acid Neutral Base

35. pOH
pOH is a measure of the concentration of hydroxide ions in a solution.
pOH = -log [OH-]

36. Example: What is the pOH of a solution where [OH-] = 1 x 10-5 M?
pOH = -log [OH-]
pOH = -log(1 x 10-5)
pOH = 5

37. How are pH and pOH related?
At every pH, the following relationships hold true:
[H3O+] • [OH-] = 1 x M
pH + pOH = 14

38. Example 1: What is the pH of a solution where [H+] = 3.4 x 10-5 M?
pH = -log [H+]
pH = -log(3.4 x 10-5 M)
pH = 4.5

39. Example 2: What is the pH of a solution where [H+] = 5.4 x 10-6 M?
pH = -log [H+]
pH = -log(5.4 x 10-6)
pH = 5.3

40. Example 3: What is the [OH-] and pOH for the solution in example #2?
[H3O+][OH-]= 1 x 10-14
(5.4 x 10-6)[OH-] = 1 x 10-14
[OH-] = 1.9 x 10-9 M
pH + pOH = 14
pOH = 14 – 5.3 = 8.7

41. Example #4 ***MUST SOLVE FOR pH and use the pH scale basic basic
Classify each solution as acidic, basic, or neutral
***MUST SOLVE FOR pH and use the pH scale
a. [H+] = 6.0 x M
b. [OH-] = 3.0 x 10-2 M
c. [H+] = 2.0 x 10-7 M
d. [OH-] = 1.0 x 10-7 M
basic
basic
acidic
neutral

42. Example #5
Which is the MOST basic from question #4? B.

43. Acids and bases: Titrations
The amount of acid or base in a solution is determined by carrying out a neutralization reaction;
an appropriate acid-base indicator (changes color in specific pH range) must be used to show when the neutralization is completed.

44. Solution with Indicator
Read a buret volume to 2 decimal places
This process is called a titration: the addition of a known amount of solution to determine the volume or concentration of another solution.
Buret
Solution with Indicator

45. Textbook page 615 Figure 19.22 a-c
End point: the point at which the indicator changes color

46. 3 Steps to do a titration (pg. 615):
(show lab in demo form…)
3 Steps to do a titration (pg. 615):
Add a measured amount of an acid of unknown concentration to a flask.
Add an appropriate indicator to the flask
Add measured amounts of a base of known concentration using a buret. Continue until the indicator shows that neutralization has occurred. This is called the end point of the titration

47. 4 steps to a titration CALC:
1) balanced equation
2) calculate the number of moles of acid or base in known solution
3) calculate the number of moles in unknown solution used during the titration
4) determine molarity of unknown solution and the pH

48. Example:
In a titration, 27.4 mL of M Ba(OH)2 is added to a 20.0 mL sample of HCl solution of unknown concentration. What is the molarity and pH of the acid solution?
Equation: (Step 1)
Ba(OH)2 + 2 HCl  BaCl2 + 2 H2O  (steps)

49. Step 2
Calculate the number of moles of known solution (Ba(OH)2)

50. Calculate moles of known solution:
Mol Ba(OH)2 =
4.22 x 10-4 mol Ba(OH)2

51. Step 3 Calculate moles of unknown solution
Use stoichiometry and the balanced equation:
Ba(OH)2 + 2 HCl  BaCl2 + 2 H2O

52. Calculate moles of unknown solution:
Mol HCl = 4.22 x 10-4 mol Ba(OH)2 x 2 mol HCl= x 10-4 mol HCl mol Ba(OH)2
Ba(OH)2 + 2 HCl  BaCl2 + 2 H2O
Use coefficients from bal. eq to get molar ratio

53. Step 4
Determine molarity and pH

54. Calculate the M and pH
Molarity = 8.44 x 10-4 mol HCl = 4.22 x 10-2 M HCl L
pH: HCl dissociates into H+ and Cl- ions
4.22 x 10-2 M HCl = 4.22 x 10-2 M H+ = 4.22 x 10-2 M Cl-
pH = -log[4.22 x 10-2 M H+] = = 1.38

55. Extra Calculation (same steps as the one we just did)
**** DO ON THE BACK OF YOUR NOTETAKERS OR CONT. IN NOTES
A 25 mL solution of H2SO4 is completely neutralized by 18 mL of 1.0 M NaOH. What is the concentration of the H2SO4 solution?
Follow steps #1-4

56. Titration Curve (pg. 615)
A graph showing how the pH changes as a function of the amount of added titrant in a titration.
Data for the graph is obtained by titrating a solution and measuring the pH after EVERY drop of added titrant.

57. Equivalence point =
The point on the curve where the moles of acid equal the moles of base
the midpoint of the steepest part of the curve is a good approximation of the equivalence point.

58. Knowledge of the equivalence point can then be used to choose a suitable indicator for a given titration;
the indicator must change color (end point) at a pH that corresponds to the equivalence point.
pg. 602 figure 19.12

59. HANDOUT: titration curve WS