1. Chapter 19 Chemistry Exam: Tuesday, May 14
Acids & Bases
Chapter 19
Chemistry
Exam: Tuesday, May 14

2. Arrhenius’ Definition of an Acid
Arrhenius defines acids and bases by the type of ions produced in solution
Acids produce H+ ions in solution
Monoprotic acids – have 1 hydrogen to ionize (HCl)
Diprotic acids – have 2 hydrogens to ionize (H2SO4)
Triprotic acids – have 3 hydrogens to ionize (H3PO4)
Produces H+ ions in solution 2

3. Arrhenius’ Definition of a Base
Arrhenius defines acids and bases by the type of ions produced in solution
Bases produce OH- ions in solution
Produces OH- ions in solution 3

4. Bronsted-Lowry Acid
Bronsted-Lowry defines acids and bases by whether they accept or donate protons (H+)
Acids lose or donate a proton (aka hydrogen ion) during a reaction 
Acid loses a proton

5. Bronsted-Lowry Base
Bronsted-Lowry defines acids and bases by whether they accept or donate protons (H+)
Bases gain or accept a proton (aka hydrogen ion) during a reaction 
Base gains a proton

6. Properties of Acids Taste sour Feel watery
Electrolytes: conduct an electrical current by forming H+ ions in solution
Corrosive
pH < 7 6

7. Properties of Bases Also referred to as alkaline Taste bitter
Feel slippery
Electrolytes: conduct an electrical current by forming OH- ions in solution
Corrosive
pH > 7 7

8. Amphoteric
Some substances can act as an acid or a base and are called amphoteric.
Ex. H2O can act as an acid and a base
Water can lose a hydrogen ion (proton)
NH3 + H2O → NH4+ + OH-
(base) (acid)
Water can gain a hydrogen ion (proton)
HCl + H2O → H3O+ + Cl-
(acid) (base) 8

9. Conjugate acid/base pairs
Conjugate acid – product that remains when a base has accepted a proton.
Conjugate base – product that remains when an acid has donated a proton.
Acids and bases are always reactants
Conjugate acids and conjugate bases are always products.
Ex NH3 + H2O → NH4+ + OH-
(base) (acid) (c.acid) (c.base)

10. Conjugates 10 10

11. Identifying acids, bases, & conjugates
NH3(g) + H3O+(aq)  NH4+(aq) + H2O(l)
CH3OH(l) + NH2-(aq)  CH3O-(aq) + NH3(g)
OH-(aq) + H3O+(aq)  H2O(l) + H2O(l)
NH2-(aq) + H2O(l)  NH3(g) + OH-(aq)
Base
Acid
C.Base
C.Acid
Acid
Base
C.Base
C.Acid
Base
Acid
C.Acid
C.Base
Base
Acid
C.Acid
C.Base

12. Naming an Acid Binary acids (H + nonmetal):
hydro ______ic acid
Ternary acids (H + polyatomic ion):
ate ions: _________ic acid
ite ions: _________ous acid
Name the following acids:
HBr
H2SO3
H3PO4
Write the formulas for the following acids:
Hydrosulfuric acid
Nitrous acid
Chromic acid 12

13. Naming a Base Formula is cation + OH- _______ hydroxide
Name the following bases:
NaOH
NH4OH
NH3
Ca(OH)2
Write the formulas for the following bases:
Magnesium hydroxide
Aluminum hydroxide 13

14. Indicators
Indicators: organic substances that change colors in an acid or a base (sometimes paper, sometimes liquids)
Acids turn…
Litmus paper – red
Phenolphthalein – clear
Universal indicator – yellow / orange / red
Bases turn…
Litmus paper – blue
Phenolphthalein –pink
Universal indicator – dark green /blue / purple 14

15. Indicators
Litmus Paper
Phenolphthalein
Universal Indicator 15 15

16. Practice Identify acid/base pairs: Name these acids/bases:
Acids produce _____ions in solution
Monoprotic acids – have ___ hydrogen to ionize (HCl)
Diprotic acids – have ___ hydrogens to ionize (H2SO4)
Triprotic acids – have ___ hydrogens to ionize (H3PO4)
Bases produce ___ ions in solution
Identify acid/base pairs:
NH2-(aq) + H2O(l)  NH3(g) + OH-(aq)
HCl + H2O → H3O+ + Cl-
Name these acids/bases:
HCl
HClO3
NaOH 16

17. The pH / pOH Scale
pH is a mathematical scale in which the concentration of hydrogen ions in a solution is expressed as a number from 0 to 14.
The pH scale is a convenient way to describe the concentration of H+ ions in acidic solutions, as well as the hydroxide ions in basic solutions (pOH)

18. Household Items & pH Scale

19. pH and pOH Calculations
The pH of a solution equals the negative logarithm of the hydrogen ion concentration.
The pOH of a solution equals the negative logarithm of the hydroxide ion concentration. pH
pOH
Stands for “power of hydrogen”
Stands for “power of hydroxide”
Shows the [H+] in a solution
Shows the [OH-] in a solution
pH = -log[H+]
pOH = -log[OH-]

20. Logarithm Tutorial/Review
In math, log is short for logarithm
A logarithm is the power to which a base, must be raised to produce a given number.
B = base, which will always be 10
X = exponent
N = argument
Ex: log 1000 = 3 can be rewritten as 103 = 1000
a) log 100 =
b) log 25 =
c) – log (0.0001) =
logBn = x also written as Bx = n 2
1.4 4

21. pH and pOH Calculations
Patterns in pH/pOH calculations
In an acidic solution, [H+] > [OH-]
In a basic solution, [H+] < [OH-]
pH + pOH = 14
Ex: If pH = 4, then pOH = 10
[H+] and [OH-] exponents always add up to -14
Ex: If [H+] = 1.0x10-3M, then [OH-] = 1.0x10-11M

22. pH and pOH Calculations
Notice the relationship between concentration and the pH/pOH value:
Toothpaste has a hydrogen ion concentration of 10–10M, so its pH is 10.
pH = -log[10-10] = 10
Pure water, which is neutral, has a pH of 7. That means its hydrogen ion concentration is 10–7M.
7 = -log[H+] ; [H+] = 10-7

23. pH & pOH Practice
In a NaOH solution the pH = 11. What is the [H+] of the solution?
Find the pH of a 0.1M HNO3 solution.
If a substance’s pH = 1, what is the pOH?
11 = -log[H+] ; [H+] = 1x10-11 M
pH = 1
pH = -log[0.1] = 1
1 + pOH = 14 ; pOH = 13 23

24. pH & pOH Practice pOH pH [OH-] [H+] 4 2 1x10-6 1x10-3 9 13 10 1x10-412
1x10-12
1x10-2 6 8
1x10-8 3 11
1x10-11 5
1x10-5
1x10-9 1
1x10-13
1x10-1 3 11
1x10-11

25. Strength of Acids and Bases
Strength is determined by how much acids and bases ionize (dissociate) in water
Strong – ionize 100% in water
Weak – only partially ionize in water
Complete dissociation - all HCl compounds have separated into H+ and Cl- ions
Partial dissociation - some HNO2 compounds have separated into H+ and NO2- ions, while some remain together 25

26. Strength of Acids and Bases
The terms weak and strong are used to compare the strengths of acids and bases
The terms dilute and concentrated are used to compare the concentration of solutions
They do not mean the same thing!
The combination of strength and concentration ultimately determines the behavior of the acid or base.

27. Strength of Acids and Bases
Strong Acids:
Only HNO3, HI, HBr, HCl, H2SO4, HClO4
Remember: NO, I Brought Claude SOme ClOthes
All other acids are weak, and remain in equilibrium
Ex: HNO2 ↔ H+ + NO2-
Strong Bases: Group I or II metals
Ex: NaOH, Mg(OH)2…etc.
All other bases are weak, and remain in equilibrium
NH3 + H2O ↔ NH4+ + OH- 27

28. Conjugate Strength Strong Acid = weak conjugate base
Strong Base = weak conjugate acid
Weak Acid = strong conjugate base
Weak Base = strong conjugate acid
Example:
HCl + NH3 → Cl NH4+
(strong acid) (weak base) (weak conjugate base) (strong conjugate acid) 28

29. Strong Acid/Base Calculations
If the acid/base is strong, concentrations of reactants will match products
Ex: What is the [H+] in a 0.2M HCl solution? HCl is a strong acid; it ionizes completely.
HCl  H Cl-
Ex: What is the [OH-] in a 0.03M NaOH solution? NaOH is a strong base; it ionizes completely.
NaOH  Na OH-
0.2M
0.0M
[H+] = 0.2M
0.03M
0.0M
[OH-] = 0.03M

30. Ionization Constants – Ka & Kb
If the acid/base is weak, concentrations must be calculated using ionization constants
Ka & Kb – acid/base dissociation constants; ratio of the concentration of the dissociated ions to the undissociated acid/base (no units)
Acid + water ↔ H+ + anion-
Base + water ↔ cation+ + OH-
Ka = [H+][anion-] & Kb = [cation+][OH-]
[acid] [base]
The closer the Ka or Kb is to 1, the greater the extent of dissociation (the stronger the acid or base)!

31. Weak Acid Calculation
Ex: A weak acid, HA, is found to be 25% ionized in water. Find the Ka for a 0.4M solution of HA.
HA  H A-
0.4M
0.0M
0.0M
M= 0.3M
0.1M
0.1M
Ka = [H+][A-]
[HA]
= [0.1][0.1]
[0.3]
= 0.033

32. Weak Base Calculation BOH  B+ + OH-
Ex: A weak base, BOH, is prepared by dissolving 1.5 moles in water to make 3 liters of solution. At equilibrium, the [OH-] is 0.01M. Find the Kb.
BOH  B OH-
0.5M
0.0M
M = 0.49M
0.01M
Kb = [B+][OH-]
[BOH]
= [0.01][0.01]
[0.49]
= 2.0x10-4

33. Additional Calculations
The Kb for NH3 is 1.8x Find the [OH-] in a 0.5M NH3 solution.
If the ionization constant is less than 5% (5x10-2), the amount of acid or base that ionizes is neglible and can be omitted.
NH H2O  NH OH-
0.0M
0.5M
0.0M x x
0.5-x
Ka = [NH4+][OH-]
[NH3]
= [x][x]
[0.5-x]
= x2
[0.5]
= 1.8x10-5
X2 = 9.0x10-6
[OH-] = X = 3.0x10-3M

34. Additional Calculations
What is the hydronium ion concentration of a 0.1M solution of formic acid (HCOOH) if the Ka = 1.77x10-4?
HCOOH + H2O  COOH- + H3O+
0.1M
0.0M
0.0M
0.1-x x x
Ka = x2
[0.1]
= 1.77x10-4
X2 = 1.77x10-5
[H3O+] = X = M

35. Additional Calculations
The Ka for a weak acid, HA is 3.0x10-5 (less than 5%). Find the pH of a 0.2M solution of HA.
HA  H A-
0.2M
0.0M
0.0M
0.2-x x x
Ka = [H+][A-]
[HA]
= x2
0.2
= 3.0x10-5
X = [H+] = 2.4x10-3M
pH = -log [H+]
= -log 2.4x10-3
= pH = 2.6

36. Additional Calculations
The pH for a 0.1M BOH (weak base) solution is Find the Kb.
XOH  X OH-
pH = so pOH =
3.5
3.5 = -log [OH-]
[OH-] = =
3.16x10-4M
x10-4M
3.16x10-4M
3.16x10-4M
Kb = [X+][OH-]
[XOH]
= (3.16x10-4)2
x10-4
= 1.0x10-6

37. Acidic and Basic Anhydrides
Anhydrides are substances without water
Acidic anhydrides (ex. CO2) could be made an acid if water is added
nonmetal oxide + water  acid
Ex. CO2 + H2O  H2CO3
Other acidic anhydrides include: SO2, NO2, N2O5
Basic anhydrides (ex. Na2O) could be made a base if water is added
metal oxide + water  base
Na2O + H2O  2NaOH
Other basic anhydrides include: SrO, Al2O3, Li2O, NO2

38. Neutralization Reactions
The reaction of an acid and a base to produce a salt and water is called a neutralization reaction.

39. Neutralization Reactions
Ex: sodium hydroxide (base) and hydrochloric acid (acid) react to form sodium chloride (salt) and water.

40. Neutralization Practice
Neutralizations are double displacement reactions!
Practice predicting products (always a salt and water):
Example #1
__Ca(OH)2 + __H3PO4 → ?
__Ca(OH)2 + __H3PO4 → __Ca3(PO4)2 + __H2O
3Ca(OH)2 + 2H3PO4 → Ca3(PO4)2 + 6H2O
Example #2
__Fe(OH)2 + __HBr → ?
__Fe(OH)2 + __HBr → __FeBr2 + __H2O
Fe(OH)2 + 2HBr → FeBr2 + 2H2O 40

41. Hydrolysis of a Salt
The ability of a salt to form an acidic, basic, or neutral solution.
Parent Acid
Parent Base
pH of Solution
Strong
Neutral
Weak
Acidic
Basic
Questionable

42. Acidic, Basic or Neutral
Hydrolysis Examples
Acid + Base  Salt + Water (neutralization)
Salt
Parent Base
Parent Acid
Acidic, Basic or Neutral
Ca3(PO4)2
FeBr2
MgI2
Al2(SO3)3
Ca(OH)2
H3PO4
Basic
strong
weak
Fe(OH)2
HBr
Acidic
weak
strong
Mg(OH)2 HI
Neutral
strong
strong
Al(OH)3
H2SO3
??????
weak
weak

43. Buffers
A buffer is a solution that resists changes in pH when moderate amounts of acids or bases are added.
Buffer solutions are prepared by using a weak acid with one of its salts or a weak base with one of its salts
Ex: a buffer solution can be prepared by using the weak base ammonia, NH3, and an ammonium salt, such as NH4Cl.

44. Why is aspirin buffered?
Aspirin is acetylsalicylic acid
When a naturally acidic or basic substance is buffered, it’s pH is balanced.
If something too acidic or basic is ingested, it can seriously harm the stomach lining
Buffered aspirin is coated in a buffering agent that will maintain the pH of the aspirin as it passes through the stomach, preventing it from dissolving until it reaches the small intestine.

45. Acid-Base Titrations Titration is used in determining
the concentration of a base by
using an acid whose concentration is known.
An indicator (phenolphthalein) is added to the standardized acid (acid of known concentration).
The unknown base is added slowly with a burette until the solution is neutralized and reaches the equivalence point (where [H+] = [OH-]).
Adding one more drop of base
changes the color of the solution to
pink. This is called the endpoint.
The equivalence point is not
necessarily the midpoint, or the point
where pH = 7

46. Titration Calculations
MaVa(#H+) = MbVb(OH-)
Note: pH at the equivalence point is not always 7.
Strong Acid + Strong Base, pH=7
Strong Acid + Weak Base, pH<7
Weak Acid + Strong Base, pH>7
Example: It took 75mL of NaOH to neutralize 50mL of 2M HCl. What is the concentration of the NaOH?
MaVa(#H+) = MbVb(#OH-)
(2M)(50mL)(1) = (x)(75mL)(1)
x = 1.33M NaOH

47. Titration Practice
It took 20mL of Ca(OH)2 to neutralize 25mL of 0.05M HCl. What is the concentration of the base?
MaVa(#H+) = MbVb(OH-)
(0.05M)(25mL)(1) = (x)(20mL)(2)
x = M