1. Unit 13
Acids and Bases

2. Properties ACIDS BASES electrolyte electrolyte sour taste bitter taste
sticky feel
slippery feel
turn litmus red
turn litmus blue
react with acids to form water and a salt (ionic compound)
react with bases to form water and a salt (ionic compound)

3. Examples
ACIDS: Most citrus fruits, tea, battery acid, vinegar, milk, soda, apples.
BASES: Common household bases include baking soda, lye, ammonia, soap, and antacids. . 3

4. Indicators
Indicators are substances that change color in the presence of an acid or a base
Indicators are made up of weak acids or weak bases
Examples of indicators include pH paper, red and blue litmus paper, and phenolphthalein

5. Acids Affect Indicators:
Blue litmus paper turns red in contact with an acid. It remains blue when in contact with a base or neutral solution. 5

6. Bases affect indicators:
Red litmus paper turns blue in contact with a base. It remains red when in contact with an acid or neutral solution.
Phenolphthalein turns pink in a base. It is colorless in an acid or neutral solution. 6

7. Definitions There are 3 definitions used to describe acids and bases:
Arrhenius
BrØnsted-Lowry
Lewis
The most traditional is Arrhenius acids and bases.

8. HCl + H2O  H+ + Cl– Definitions – + Arrhenius - In aqueous solution…
Acids form hydrogen ions (H+)
HCl + H2O  H+ + Cl–
Also called hydronium ions (H3O+) H Cl O – +
acid

9. NH3 + H2O  NH4+ + OH- Definitions – +
Arrhenius - In aqueous solution…
Bases form hydroxide ions (OH-)
NH3 + H2O  NH4+ + OH- H N O – +
base

10. Another common way to refer to hydrogen ions is to call them “protons”
Definitions
Another common way to refer to hydrogen ions is to call them “protons”
Brønsted-Lowry
Acids are proton (H+) donors.
Bases are proton (H+) acceptors.
HCl + H2O  Cl– + H3O+
acid
conjugate base
base
conjugate acid
Conjugate acid – particle formed when a base gains a H+
Conjugate base – particle that remains when an acid has donated a H+ .

11. Definitions Lewis Lewis base Lewis acid
Acids are electron pair acceptors.
Bases are electron pair donors.
Lewis base
Lewis acid

12. White Board Questions
When you wafted a substance your nose burned. Would this substance be an acid or a base?
A hydrogen ion (H+) can also be called a _________ or ____________.
Arrhenius acids are compounds that break up in water to give off _____________.
What color litmus paper would you use to test an acid? What color will it turn?
5. If your food tastes bitter, which do you think it could possibly be an acid or a base?
ACID
Proton
H3O+ H+
Blue turns red
BASE 12

13. White Board Questions
6. A BrØnsted-Lowry base _________ hydrogen ions. 7. Phenolphthalein turns pink when it comes in contact with a(n) _________. 8. Which of the scientists defined the typical acid? 9. If you are eating and it has a sour taste, would that be an acid or a base? 10. If a piece of red litmus paper turns blue than it is a(n) ___________.
accepts
base
Arrhenius
acid
base 13

14. Naming Acids Binary acids Contains 2 different elements: H and another
Always has “hydro-” prefix
Root of other element’s name
Ending “-ic”

15. Examples of Binary Acids
HI is hydroiodic acid
H2S is hydrosulfuric acid
HBr is hydrobromic acid
HCl is hydrochloric acid

16. Naming Acids Ternary Acids - Oxyacids
Contains 3 different elements: H, O, and another
No prefix
Name of polyatomic ion (p. 147)
Ending “–ic” for polyatomic ion ending in “-ate” and “–ous” for ion ending in “-ite”

17. Examples of Ternary Acids
ClO3 is chlorate so HClO3 is chloric acid
PO4 is phosphate
so H3PO4 is phosphoric acid
PO3 is phosphite
so H3PO3 is phosphorous acid
NO2 is nitrite HNO2 is nitrous acid
NO3 is nitrate HNO3 is nitric acid

18. Naming Acids cont.
HC2H3O2 or CH3COOH
Name is acetic acid
Common name = vinegar

19. Practice Naming Acids H2SO3 Sulfurous acid HF Hydrofluoric acid H2Se
Hydroselenic acid
Perchloric acid
HClO4
Carbonic acid
H2CO3
Hydrobromic acid
HBr

20. Ion Product of Water H2O H+ + OH-
Self- ionization of water – the simple dissociation of water
H2O H OH-
Concentration of ea. ion in pure water:
[H+] = 1.0 x 10-7M [OH-] = 1.0 x 10-7M
Ion-product constant for water (Kw), Where Kw = 1.0 x 10-14
Kw = [H+] [OH-]
Acid [H+] > [OH-]
Base [H+] < [OH-]
Neutral [H+] = [OH-] 20

21. Calculating [H+] and [OH-]
reverse the pH equation
The pH of a solution is 8. Find the [H+] and [OH-] and determine whether it is acidic, basic, or neutral.
basic
[H+] = 1 x 10-pH and [OH-] = 1 x 10-pOH
[H+] = 1 x 10-8 M
[OH-] = 1 x 10-(14-8) M = 1 x 10-6 M

22. Examples
1. If the [H+] in a solution is 1.0 x 10-5M, is the solution acidic, basic or neutral?
1.0 x 10-5 M
What is the concentration of the [OH-]?
Use the ion-product constant for water (Kw):
Kw = [H+] [OH-]
1.0 x = [1.0 x 10-5] [OH-]
1.0 x = [OH-]
1.0 x 10-5
1.0 x 10-(14-5)
pH 5 = acidic
1.0 x 10-9 M 22

23. Examples
2. If the pH is 9, what is the concentration of the hydroxide ion?
Kw = [H+] [OH-]
1.0 x M = [1.0 x 10-9M] [OH-]
1.0 x 10-5 M = [OH-]
14 = pH + pOH
14 = 9 + pOH
5 = pOH
3. If the pOH is 4, what is the concentration of the hydrogen ion?
Kw = [H+] [OH-]
1.0 x M = [H+] [1.0 x 10-4 M]
1.0 x M = [H+]
14 = pH + pOH
14 = pH + 4
10 = pH 23

24. Examples Acidic since pH is 4
4. A solution has a pH of 4. Calculate the pOH, [H+] and [OH-]. Is it acidic, basic, or neutral?
14= pH + pOH
14= 4 + pOH
10= pOH
Acidic since pH is 4

25. Practice Problems: Classify each solution as acidic, basic or neutral.
1. [H+] = 1.0 x M
2. [H+] = 0.001M
3. [OH-] = 1.0 x 10-7 M
4. [OH-] = 1.0 x 10-4 M
Basic pH 10
1.0 x 10-3 acid pH 3
Neutral
14=pH+4 base pH 10 25

26. [OH-] pOH pH [H+] 1 x 10-14 14 1 x 100 1 x 10-13 13 1 1 x 10-1
1 x 100
1 x 10-13 13 1
1 x 10-1
1 x 10-12 12 2
1 x 10-2
1 x 10-11 11 3
1 x 10-3
1 x 10-10 10
Increasing acidity 4
1 x 10-4
1 x 10-9 9 5
1 x 10-5
1 x 10-8 8 6
1 x 10-6
1 x 10-7 7
Neutral
Increasing basicity 26 26

27. White Board Practice Fill in the chart.
[OH-]
pOH pH
[H+] 8
1x 10-12 10
1 x 10-3
1.0 X 10 -8 6
1.0 X 10 -6 2 12
1.0 X 10 -2
1.0 X 10 -4 4
1.0 X 11 3
1.0 X

28. Fill in the chart. [OH-] pOH pH [H+] 8 1x 10-12 10 1 x 10-3 5 1 × 10-16
1.0 X 10 -6 2 12
1.0 X 10 -2
1.0 X 10 -4 4
1.0 X 11 3
1.0 X 9
1.0 X 10 -5
1.0 X 10 -9 13 1
1.0 X

29. Strength or Concentration- +
Strong Acid/Base
Ionize completely in water
strong electrolyte
Acids
HCl
HNO3
H2SO4
HBr HI
HClO4
Bases
NaOH
KOH
Ca(OH)2
Ba(OH)2

30. Strength or Concentration
Weak Acid/Base
ionize partially in water
weak electrolyte - +
Acids HF
CH3COOH
H3PO4
H2CO3
HCN
Base
NH3

31. Strength or Concentration
How strong or weak an acid or base is, depends on its degree of ionization. - + - +

32. pouvoir hydrogène (Fr.)
pH Scale 14 7
INCREASING
ACIDITY
INCREASING
BASICITY
NEUTRAL
pH is the negative logarithm of the hydrogen ion concentration
pH = -log[H+]
pouvoir hydrogène (Fr.)
“hydrogen power”

33. The pH Scale 33

34. pH of Common Substances
pH Scale
pH of Common Substances

35. pH = -log[H+] pOH = -log[OH-] pH + pOH = 14
pH formulas
pH = -log[H+]
pOH = -log[OH-]
pH + pOH = 14

36. Neutralization Chemical reaction between an acid and a base.
Products are a salt (ionic compound) and water.

37. ACID + BASE  SALT + WATER
Neutralization
ACID + BASE  SALT + WATER
HCl + NaOH  NaCl + H2O
strong
strong
neutral
HC2H3O2 + NaOH  NaC2H3O2 + H2O
weak
strong
basic
Salts can be neutral, acidic, or basic.
Neutralization does not mean pH = 7.

38. Titration
standard solution
unknown solution
Titration
Analytical method in which a standard solution is used to determine the concentration of an unknown solution.

39. Titration cont. Equivalence point (endpoint)
Point at which equal amounts of H+ and OH- have been added.
Determined by…
indicator color change
dramatic change in pH

40. moles H+ = moles OH- MV n = MV n Titration formula M: Molarity
V: volume
n: # of H+ ions in the acid or OH- ions in the base

41. Titration example
42.5 mL of 1.3M KOH are required to neutralize 50.0 mL of H2SO4. Find the molarity of H2SO4.
H2SO4
M = ?
V = 50.0 mL
n = 2
KOH
M = 1.3M
V = 42.5 mL
n = 1
MVn = MVn
M(50.0mL)(2)=(1.3M)(42.5mL)(1)
M= 55.25
100
M = 0.55M H2SO4

42. Review of Acid and Base Definitions
Arrhenius
Most specific/exclusive definition
Created by Svante Arrhenius, Swedish
Acid: compound that creates H+ in an aqueous solution
HNO3  H+ + NO3-
Base: compound that creates OH- in an aqueous solution
NaOH  Na+ + OH-

43. Review of Acid and Base Definitions
Bronsted-Lowry
More general definition than Arrhenius definition
Most commonly used definition
Created by 2 scientists around the same time (1923)
Acid: Molecule or ion that is a proton (H+) donor
HCl + H2O  H3O+ + Cl-
Base: Molecule or ion that is a proton (H+) acceptor
NH3 + H2O  NH4+ + OH-

44. Review of Acid and Base Definitions
Lewis
Most general definition
Defined by electrons and bonding rather than H+
Created by the same scientist who electron-dot diagrams are named after
Acid: atom, ion, or molecule that accepts an electron pair to form a covalent bond
NH3 + Ag+  [Ag(NH3)2]+
Base: atom, ion, or molecule that donates an electron pair to form a covalent bond
BF3 + F-  BF4-