1. Definitions and Properties
Acids and Bases
Definitions and Properties

3. Acidic or Basic?

4. Common Acids and Bases Strong acids/bases – 100% dissociation into ions Weak acids/bases – partial dissociation, both ions and molecules
Strong Acids (strong electrolytes)
Strong Bases (strong electrolytes)
NaOH sodium hydroxide
KOH potassium hydroxide
Ca(OH) calcium hydroxide
HCl hydrochloric acid
HNO3 nitric acid
HClO4 perchloric acid
H2SO4 sulfuric acid
Weak Acids (weak electrolytes)
CH3COOH acetic acid
H2CO3 carbonic
Weak Base (weak electrolyte)
NH ammonia
NH3 + H2O  NH4OH

5. In the Context of Acids and Bases, What is a Proton?
It is still a singly charged nuclear particle.
But, typically it refers to H+, a hydrogen atom stripped of its electron and, thus, in its ionic form.
One adds a proton to a base: add H and one unit of positive charge.
One removes a proton from an acid: subtract H and one unit of positive charge.

6. 100% dissociation of HA HA H+ Strong Acid A-
Would the solution be conductive?

7. At any one time, only a fraction of the molecules are dissociated.
HA  H A- HA H+
Weak Acid A-
At any one time, only a fraction of the molecules are dissociated.

8. pH Scale Measures Acidity/Basicity
ACID
BASE
NEUTRAL
Steps on the pH scale represents a 10X change.
It’s a logarithmic scale..
pH 5 vs. pH (10X more acidic) pH 3 vs. pH (100X different) pH 8 vs. pH (100,000X different)

9. pH = -log [H+] [H+] is the concentration of the H+ ion in solution
Acid
Base
[H+] =
[OH-]
[H+]
Acidic
Basic
[OH-]
Neutral
pH = 7

10. pH of Common Substances
vinegar
2.8
water (pure)
7.0
soil
5.5
gastric
juice
1.6
carbonated
beverage
3.0
drinking water
7.2
bread
5.5
potato
5.8
blood
7.4
1.0 M
NaOH
(lye)
14.0
orange
3.5
1.0 M
HCl
milk of
magnesia
10.5
apple juice
3.8
urine
6.0
detergents
bile
8.0
lemon
juice
2.2
tomato
4.2
milk
6.4
ammonia
11.0
seawater
8.5
coffee
5.0
bleach
12.0 1 2 3 4 5 6 7 8 9 10 11 12 13 14
acidic
neutral
basic
[H+] = [OH-]

11. Common Acids Sulfuric Acid H2SO4 Nitric Acid HNO3
Phosphoric Acid H3PO4
Hydrochloric Acid HCl
Acetic Acid CH3COOH
Carbonic Acid H2CO3
Battery acid
Used to make fertilizers and explosives
Food flavoring
Stomach acid
Vinegar
Carbonated water

12. Common Bases NH4OH NH4+ + OH- ammonium hydroxide
Name Formula Common Name
Sodium hydroxide NaOH lye or caustic soda (Drano)
Potassium hydroxide KOH lye or caustic potash
Magnesium hydroxide Mg(OH)2 milk of magnesia
Calcium hydroxide Ca(OH)2 slaked lime
Ammonia water NH3 H2O household ammonia
NH4OH
NH OH-
ammonium hydroxide

13. Acid/Base Definitions
Lewis Definitions
A Lewis acid is a substance than can accept (and share) an electron pair.
A Lewis base is a substance than can donate (and share) an electron pair.
Lewis Acid
Brønsted-Lowry Definitions
A Brønsted-Lowry acid is a proton donor; it donates a hydrogen ion, H+.
A Brønsted-Lowry base is a proton acceptor; it accepts a hydrogen ion, H+.
Brønsted- Lowry A
Arrhenius Acids and Bases
Acids release hydrogen ions in water.
Bases release hydroxide ions in water.
An acid is a substance that produces hydronium ions, H3O+, when dissolved in water.

14. Electron-pair acceptor
Acid – Base Systems
Type
Acid
Base
Arrhenius
H+ or H3O + producer
OH - producer
Brønsted-
Lowry
Proton (H +) donor
Proton (H +) acceptor
Lewis
Electron-pair acceptor
Electron-pair donor

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

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

17. HCl + H2O  Cl– + H3O+ Definitions acid conjugate base base
Brønsted-Lowry
Acids are proton (H+) donors.
Bases are proton (H+) acceptors.
HCl + H2O  Cl– + H3O+
acid
conjugate base
base
conjugate acid

18. Water is Amphoteric Amphoteric Substances:
Substances which can act as either proton donors (acids) or proton acceptors (bases) depending on what substances are present.
HCl H2O H3O Cl-
acid base CA CB
base acid CA CB
NH H2O NH OH-

19. Definitions
H2O + HNO3  H3O+ + NO3– B A CA CB
Base
Acid H O N H O

20. Definitions
NH3 + H2O  NH4+ + OH- B A CA CB
Base
Acid H H O H N H

21. HF H3PO4 H3O+ F - H2PO4- H2O Conjugate Bases for Acids
Give the conjugate base for each of the following: HF
H3PO4
H3O+
F -
H2PO4-
H2O
Polyprotic - an acid with more than one H+

22. Conjugate Acids for Bases
Give the conjugate acid for each of the following:
Br -
HSO4-
CO32-
HBr
H2SO4
HCO3-
Courtesy Christy Johannesson

23. ACID + BASE  SALT + WATER
Neutralization
Neutralization is a chemical reaction between an acid and a base to produce a salt (ionic compound) and water.
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.

24. Examples of everyday uses of neutralization
Shampoo (a mild alkali) and conditioner (a mild acid)
Toothpaste (alkaline) neutralizes acids produced by bacteria
Wasp sting (alkaline) and vinegar to relieve pain
Acidic soil neutralize by adding calcium hydroxide (lime)

25. Reactions that Produce Salt (Double Replacement Reactions)
acid base salt + water
H3PO4 +
NH4OH
(NH4)3PO4 +
H2O
phosphoric acid
ammonium hydroxide
ammonium phosphate
water
HNO3
Mg(OH)2
Mg(NO3)2
H2O
nitric acid
magnesium hydroxide
magnesium nitrate
H2CO3
KOH
K2CO3
H2O
carbonic acid
potassium hydroxide
potassium carbonate
HC2H3O2
Al(OH)3
Al(C2H3O2)3
H2O
acetic acid
aluminum hydroxide
aluminum acetate
HClO4
Ba(OH)2
Ba(ClO4)2
H2O
perchloric acid
barium hydroxide
barium perchlorate

26. pH Indicators and Titration
Acids and Bases
pH Indicators and Titration pH pH

27. pH of Common Substances
vinegar
2.8
water (pure)
7.0
soil
5.5
gastric
juice
1.6
carbonated
beverage
3.0
drinking water
7.2
bread
5.5
potato
5.8
blood
7.4
1.0 M
NaOH
(lye)
14.0
orange
3.5
1.0 M
HCl
milk of
magnesia
10.5
apple juice
3.8
urine
6.0
detergents
bile
8.0
lemon
juice
2.2
tomato
4.2
milk
6.4
ammonia
11.0
seawater
8.5
coffee
5.0
bleach
12.0 1 2 3 4 5 6 7 8 9 10 11 12 13 14
acidic
neutral
basic
[H+] = [OH-]

28. Common Acid-Base pH Indicators
Range and Color Changes of Some
Common Acid-Base pH Indicators
pH Scale 1 2 3 4 5 6 7 8 9 10 11 12 13 14
Indicators
Methyl orange red – yellow
Methyl red red yellow
Bromthymol blue yellow blue
Neutral red red yellow
Phenolphthalein colorless red colorless beyond 13.0
Bromthymol blue indicator would be used in titrating a strong acid with a strong base.
Phenolphthalein indicator would be used in titrating a weak acid with a strong base.
Methyl orange indicator would be used in titrating a strong acid with a weak base.

29. Indicator:
A solution that changes color in response to changes in a solution’s H+ (acidity) and OH- (basicity) concentrations.

30. Desired Features of pH Indicators
pH paper pH pH
Detection limit
Low deflection
High sensitivity
High selectivity
Wide dynamic range
Simple to use
Cost-effective

31. pH

32. INDICATOR COLORS IN TITRATION
Acid color
Transition color
Base color
STRONG ACID – STRONG BASE
Litmus pH
Bromthymol blue

33. INDICATOR COLORS IN TITRATION
Indicator
Acid color
Transition color
Base color
Phenolphthalein
Phenol red
WEAK ACID – STRONG BASE pH

34. INDICATOR COLORS IN TITRATION
Indicator
Acid color
Transition color
Base color
Methyl orange
Bromphenol blue
STRONG ACID – WEAK BASE pH

35. Common pH Indicators
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 520

36. Edible Acid-Base Indicators
COLOR CHANGES AS A FUNCTION OF pH
INDICATOR pH
RED APPLE SKIN
BEETS
BLUEBERRIES
RED CABBAGE
CHERRIES
GRAPE JUICE
RED ONION
YELLOW ONION
PEACH SKIN
PEAR SKIN
PLUM SKIN
RADISH SKIN
RHUBARB SKIN
TOMATO
TURNIP SKIN *
*YELLOW at pH 12 and above

37. Red Cabbage Indicator

38. Neutralization Reaction
A chemical reaction between an acid and a base that destroys the distinctive properties of both.
A neutralization reaction produces water and a salt.
Before:
After:

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

40. Acid-Base Titration Buret:
Titration glassware for delivering measured drops of solution. The graduations on this piece of equipment start at 0.0 mL at the top and increase as one goes down towards the flow control stopcock.
Buret

41. Titration Equivalence point (endpoint)
Point at which equal amounts of H3O+ and OH- have been added. Often confused with the “endpoint.”
Determined by…
indicator color change
dramatic change in pH

42. moles H3O+ = moles OH- MV n = MV n Titration M: Molarity V: volume
n: # of H+ ions in the acid or OH- ions in the base

43. Titration
42.5 mL of 1.3M KOH are required to neutralize 50.0 mL of H2SO4. Find the molarity of H2SO4.
H3O+
M = ?
V = 50.0 mL
n = 2
OH-
M = 1.3M
V = 42.5 mL
n = 1
MV# = MV#
M(50.0mL)(2)
=(1.3M)(42.5mL)(1)
M = 0.55M H2SO4

44. Titration Curve

45. Titration of a Strong Acid With a Strong Base
(20.00 mL of M HCl by M NaOH)
14.0
12.0
Color change
alizarin yellow R
10.0
Color change
phenolpthalein
8.0
equivalence
point
Color change
bromthymol blue pH
6.0
Color change
bromphenol blue
4.0
Color change
methyl violet
2.0
0.0
0.0
10.0
20.0
30.0
Volume of M NaOH added
(mL)

46. Titration of a Weak Acid With a Strong Base
NaOH added
(mL) pH
Titration Data
14.0
12.0
10.0
equivalence point
8.0 pH
6.0
4.0
2.0
0.0
0.0
10.0
20.0
30.0
40.0
Volume of M NaOH added
(mL)
Phenolphthalein is best indicator: pH change

47. Titration of a Weak Base With a Strong Acid
HCl added
(mL) pH
Titration Data
14.0
12.0
10.0
8.0 pH
6.0
equivalence point
4.0
2.0
0.0
0.0
10.0
20.0
30.0
40.0
50.0
Volume of M HCl added
(mL)
Methyl orange is best indicator: pH change 3.1 – 4.4