1. AP Chemistry
Acids & Bases, Ksp

2. Acids and bases "ACID"--Latin word acidus, meaning sour. (lemon)
"ALKALI"--Arabic word for the ashes that come from burning certain plants;
water solutions feel slippery and taste bitter. (soap)

3. Acid-base theories Arrhenius definition
Acid – donates a hydrogen ion (H+) in water
Base – donates a hydroxide ion in water (OH−)
This theory is limited to substances with those “parts”. Ammonia is a MAJOR exception!
Acid examples: HCl, H2SO4
Base examples: NaOH, KOH, Ba(OH)2

4. What is the Bronsted Lowry Theory | The Chemistry Journey

5. Acid-base theories Bronsted-Lowry Definition
Acid – donates a proton (H+) in water
Base – accepts a proton in water
This theory is better; it explains ammonia as a base! This is the main theory that we will use for our acid/base discussion.

6. Acid-base theories Lewis Definition Acid – accepts an electron pair
Base – donates an electron pair
This terminology is not as widely used as it used to be because it is too general. Instead, we focus on coordination complexes, which are products of Lewis Acid-Base reactions.

7. THE BRONSTED-LOWRY CONCEPT OF ACIDS AND BASES
Using this theory, you should be able to write weak acid/base dissociation equations and identify acid, base, conjugate acid and conjugate base.
conjugate acid-base pair – A pair of compounds that differ by the presence of one H+ unit

8. Neutral compound as an acid:
HNO3 + H2O  H3O+ + NO3 −
Conj.
Acid
Conj.
Base
ACID
BASE

9. cation as an acid: NH4+ + H2O  H3O+ + NH3 Conj. Acid Conj. Base ACID

10. Anion as an acid: H2PO4− + H2O  H3O+ + HPO42− ACID BASE Conj. Acid
In each of the acid examples notice the formation of H3O+ – this species is named the “hydronium ion”
It lets you know that the solution is acidic!

12. Basic solutions
Label the acid, base, conjugate acid, and conjugate base.
NH3 + H2O  NH4+ + OH-
CO32- + H2O  HCO3- + OH-
PO43- + H2O  OH- + HPO42-
Notice the formation of OH- in each of the alkaline examples. This species is named the hydroxide ion. It lets you know the resulting solution is basic!

14. Conjugate Acids and Bases | The Chemistry Journey

15. Amphoteric molecules
Amphoteric --molecules or ions that can behave as EITHER acids or bases; water, anions of weak acids (look at the examples above—sometimes water was an acid, sometimes it acted as a base)

16. Neutralization reaction
Video containing information about water as an amphoteric molecule.

17. Bell work – 11/29/16
For the following reactions, identify the acid, base, conjugate acid, and conjugate base.
C2H5OH + H2O  H3O+ + C2H5O-
ACID BASE CA CB
HPO42- + NH4+  NH3 + H2PO4-
BASE ACID CB CA

18. pH Scale

19. pH Scale 7 14

20. Why does [H+] = [OH-] at pH 7?

21. pH and pOH pH
[H+] (M)
[OH-] (M)
pOH 7 6 3 12

22. pH and pOH What is the pattern between pH and [H+]?
What is the relationship between [H+] and [OH-]?
How is pH related to pOH?

23. Calculating pH

24. Calculating ph pH = – log[H+] pOH = – log[OH–]
The pH of a solution is 3.75, what is the Hydrogen ion concentration, [H+]?

25. Examples
What is the pH of a solution with a hydrogen ion concentration of 2.0 x 10-11?
What is the pOH of a solution with a hydroxide ion concentration of 1.8 x 10-4?
The pH of human blood was measured to be 7.41 at 25°C. Calculate the pOH, [H+], and [OH-] for the sample.
10.70
3.74
pOH = 6.59, [H+] = 3.9 x 10-8 M, [OH-] = 2.6 x 10-7 M

26. Acid and base quiz 1. OH- + HPO42-  H2O + PO43-
For questions 1 and 2 identify the acid, base, conjugate acid, and conjugate base.
1. OH- + HPO42-  H2O + PO43-
2. HF + NH3  NH4+ + F-
For questions 3-7 Name the acid or base.
H3PO HNO2
HBr LiOH
Mg(OH)2
Given a pH of 9, find the [H+], pOH, and [OH-]. Is the solution acidic or basic? (show all work)
If [H+] = 2.5 x 10-5, find the pH, pOH, and [OH-]. Is the solution acidic or basic? (show all work)

27. After quiz……
Predict the products for the following neutralization reactions:
NaOH + HCl 
HBr + Mg(OH)2 

28. Concentrated vs. Strong
“Concentrated” – refers to the amount dissolved in solution.
“Strong” – refers to the fraction of molecules that ionize.  
For example, if you put a lot of ammonia into a little water, you will create a highly concentrated solution. However, since only 0.5% of ammonia molecules ionize in water, this basic solution will not be very strong.

29. Acid-Base Titration
Uses a neutralization reaction to determine the concentration of an acid or base.
Standard Solution: the reactant that has a known molarity
Endpoint: the point at which the unknown has been neutralized.

30. Make sure you finish with correct units!!
Titration Examples
Example #1) .08 L of 0.100M NaOH is used to neutralize .02 L of HCl. What is the molarity of HCl? NaOH + HCl  NaCl + H2O
.08 L NaOH x .100 mols NaOH x 1 mol HCl x = M
1 L NaOH mol NaOH L HCl
Molarity of NaOH
Molar Ratio
Make sure you finish with correct units!!

31. Titration Examples (cont’d)
Example #2) A 0.1M Mg(OH)2 solution was used to titrate an HBr solution of unknown concentration. At the endpoint, 21.0 mL of Mg(OH)2 solution had neutralized 10.0 mL of HBr. What is the molarity of the HBr solution?

32. Titration Practice
What is the molarity of an Al(OH)3 solution if 30.0 mL of the solution is neutralized by 26.4 mL of a 0.25 M HBr solution?
A 0.3 M Ca(OH)2 solution was used to titrate an HCl solution of unknown concentration. At the endpoint, 35.0 mL of Ca(OH)2 solution had neutralized 10.0 mL of HCl. What is the molarity of the HCl solution?
When 34.2 mL of a 1.02 M NaOH solution is added from a buret to mL of a phosphoric acid solution that contains phenolphthalein, the solution changes from colorless to pink.  What is the molarity of the phosphoric acid?

35. HA(aq) + H2O(l) ⇄ H3O+(aq) + A– (aq)
Acid strength
A strong acid completely ionizes in solution
A weak acid only partially ionizes.
The strength of an acid depends on the equilibrium:
HA(aq) + H2O(l) ⇄ H3O+(aq) + A– (aq)
If the equilibrium lies far to the left, the acid is weak (only a small percentage of the acid molecules ionize)
If the equilibrium lies far to the right, the acid is strong (completely ionizes)

36. Strong acids HCl HNO3 HBr HClO4 HI H2SO4
Hydrochloric acid
Nitric acid
Hydrobromic acid
Perchloric acid
Hydriodic acid
Sulfuric acid
HCl
HNO3
HBr
HClO4 HI
H2SO4
All have WEAK conjugate bases!

37. Strong acids An HCl solution contains virtually no intact HCl.
Single arrow indicates complete ionization.
HCl(aq) + H2O(l)  H3O+(aq) + Cl– (aq)
An HCl solution contains virtually no intact HCl.
The HCl has essentially all ionized to form H3O+(aq) and Cl–(aq).

38. Describing concentrations of strong acids
A 1.0 M HCl solution has an H3O+ concentration of 1.0 M.
If [HCl] = 1.0 M, then [H3O+] = 1.0 M

39. Weak acids
Double arrow indicates partial ionization.
HF(aq) + H2O(l) ⇄ H3O+(aq) + F – (aq)
An HF solution contains a large number of intact (or un-ionized) HF molecules.
It also contains some H3O+(aq) and F –(aq).

40. Describing concentrations of weak acids
A 1.0 M HF solution has an [H3O+] that is much less than 1.0 M because only some of the HF molecules ionize to form H3O+.
[H3O+] << [HF]

41. Some weak acids HF H2SO3 HC2H3O2 H2CO3 HCHO2 H3PO4
Hydrofluoric acid
Sulfurous acid
Acetic acid
Carbonic acid
Formic acid
Phosphoric acid HF
H2SO3
HC2H3O2
H2CO3
HCHO2
H3PO4
All have STRONG conjugate bases!

42. Chemistry 12.4 Strength of Acids and Bases & Intro to Ka

43. ACID IONIZATION CONSTANT (Ka)
A way to quantify the relative strength of a weak acid
The equilibrium constant of the weak acid
HA(aq) ⇄ H+(aq) + A–(aq)
𝐾 𝑎 = [ H 3 O + ][ A − ] [HA]
= [ H + ][ A − ] [HA]

44. Ka
The smaller the constant, the less the acid ionizes… ∴ the weaker the acid
Hydrocyanic acid (HCN) Ka = 4.9 x 10-10
Nitrous acid (HNO2) Ka = 4.6 x 10-2
Acetic acid (HC2H3O2) Ka = 1.8 x 10-5
Hydrofluoric acid (HF) Ka = 3.5 x 10-4
Chlorous acid (HClO2) Ka = 1.1 x 10-2

45. Base ionization constant
The equilibrium expression for a base is known as Kb.
This expression is used to quantify the relative strength of a weak base.
B + H2O  HB+ + OH-
Kb = [HB+] [OH-]
[B]

46. Autoionization of water (Video)

47. Autoionization of water
Fredrich Kohlraush, around 1900 found that no matter how pure water is, it still conducts a minute amount of electric current.
This proves that water self-ionizes.
Remember - water is amophoteric!
Only about 2 in a billion water molecules are ionized at any instant!!

48. Autoionization of water

49. H+/OH- Equilibrium [H+] [OH-] [H+] [OH-]
The equilibrium expression used here is referred to as the autoionization constant for water, Kw.
Kw =
Kw =
[H+] [OH-]
[H+] [OH-]

50. Acidic or basic or neutral?
It is important to recognize the meaning of Kw. In an aqueous solution at 25°C, no matter what it contains, [H+] x [OH-] = 1.0 x
There are three possible situations:
A neutral solution, where [H+] = [OH-].
An acidic solution, where [H+] > [OH-].
A basic solution, where [H+] < [OH-].

51. Other temperatures…
If at a temperature OTHER THAN 25°C, the Kw will be a different value.
Remember that the numerical value of Keq is dependent on temperature!

52. Example 1 𝑲 𝒘 = H + OH − =𝟏.𝟎× 𝟏𝟎 −𝟏𝟒 (since it’s at 25°C)
Calculate [H+] or [OH-] as required for each of the following solutions at 25°C, and state whether the solution is neutral, acidic, or basic.
1.0 x 10-5 M OH-
1.0 x 10-7 M OH-
𝟏.𝟎× 𝟏𝟎 −𝟏𝟒 = H + (𝟏.𝟎× 𝟏𝟎 −𝟓 )
[OH-] > [H+]
∴ basic
H + = 𝟏.𝟎× 𝟏𝟎 −𝟏𝟒 𝟏.𝟎× 𝟏𝟎 −𝟓 = 𝟏.𝟎× 𝟏𝟎 −𝟗
[OH-] = [H+]
∴ neutral
𝟏.𝟎× 𝟏𝟎 −𝟏𝟒 = H + (𝟏.𝟎× 𝟏𝟎 −𝟕 )
H + = 𝟏.𝟎× 𝟏𝟎 −𝟏𝟒 𝟏.𝟎× 𝟏𝟎 −𝟕 = 𝟏.𝟎× 𝟏𝟎 −𝟕

53. Example 1 𝑲 𝒘 = H + OH − =𝟏.𝟎× 𝟏𝟎 −𝟏𝟒 (since it’s at 25°C)
Calculate [H+] or [OH-] as required for each of the following solutions at 25°C, and state whether the solution is neutral, acidic, or basic.
10.0 M H+
𝟏.𝟎× 𝟏𝟎 −𝟏𝟒 =(𝟏𝟎.𝟎) OH −
[H+] > [OH-]
∴ acidic
OH − = 𝟏.𝟎× 𝟏𝟎 −𝟏𝟒 𝟏𝟎.𝟎 = 𝟏.𝟎× 𝟏𝟎 −𝟏𝟓

54. At 10°C, the autoionization of water equals 2. 9 x 10-15
At 10°C, the autoionization of water equals 2.9 x Calculate [H+] and [OH-] as necessary under each of the following conditions:
Calculate [H+] if [OH-] = 9.3 x 10-4 M. Is the solution acidic or basic?
Calculate [H+] and [OH-] for a neutral solution.
Calculate [OH-] if [H+] = 6.7 x M. Is the solution acidic or basic?

55. Calculating the ph of a weak acid
Write the balanced equation
Write the equilibrium expression
Use the ICE Box method
I – Initial
C – Change
E - Equilibrium

56. Calculating the ph of a weak acid
Ex. 1 Calculate the pH of a 1.00 x 10-4 M solution of acetic acid. (Ka = 1.8 x 10-5) HC2H3O2 ⇌ H+ + C2H3O2 I 1.00 x C - x + x + x E 1.00 x 10-4 –x x x Ka = x2 we can ignore the “- x” because 1.00 x 10-4 –x it is so small. 1.8 x 10-5 = x x x 10-5 = x pH = -log (4.24 x 10-5 ) pH = 4.37
Ka = [H+][C2H3O2]
[HC2H3O2 ]

57. Calculating the ph of a weak acid
The hypochlorite ion (OCl-) is a strong oxidizing agent found in household bleaches and disinfectants. It also has a relatively high affinity for protons (it is a much stronger base than Cl- for example) and forms the weakly acidic hypochlorous acid (HOCl, Ka = 3.5 x 10-8).
Calculate the pH of a M aqueous solution of hypochlorous acid.

58. Calculating the ph of a weak acid

59. Calculating the ph of a weak base
The same steps are used to calculate weak bases and weak acids. (Make and ICE box).
Remember! You are now calculating pOH and [OH-]!!

60. Calculating the ph of a weak base
Calculate the pH for a 15 M solution of NH3
(Kb = 1.8 x 10-5)

61. Percent ionization of a weak acid
Concentration of ionized acid (at equilibrium)
% ionization= [ H + ] eq [HA] 0 ×𝟏𝟎𝟎
Concentration of acid (initially)

62. Examples of percent ionization
What is the percent ionization in a M solution nitrous acid if the [H3O+] is M?
If a 2.5 M solution of nitrous acid is 1.4% ionized, what is the hydronium concentration in the solution? What is the pH of the solution?

63. Examples of percent ionization
What is the percent ionization in a M solution nitrous acid if the [H3O+] is M?
If a 2.5 M solution of nitrous acid is 1.4% ionized, what is the hydronium concentration in the solution? What is the pH of the solution?
% ionization DECREASES with increasing concentration of acid

64. Percent ionization of a weak acid
Concentration of ionized acid (at equilibrium)
% ionization= [ H + ] eq [HA] 0 ×𝟏𝟎𝟎
Concentration of acid (initially)
Determine equilibrium concentrations for all species in the ionization reaction (Use ICE table if needed).
Use Ka and the equilibrium constant expression to find [H+].
Use [H+]eq and the initial acid concentration, [HA]0, to solve for percent ionization.

65. [HNO2]eq = __________________ [H+]eq = _________ [NO2-]eq = _________
Determine equilibrium concentrations for all species in the ionization reaction (Use ICE table if needed).
Use Ka and the equilibrium constant expression to find [H+].
Use [H+]eq and the initial acid concentration, [HA]0, to solve for percent ionization.
Find the percent ionization of a 2.5 M HNO2 solution. The Ka of nitrous acid is 4.6 x 10-4.
HNO2 ⇄ H+ + NO2-
[HNO2]eq = __________________
[H+]eq = _________
[NO2-]eq = _________
2.5 – x ≈ 2.5 x x

66. Determine equilibrium concentrations for all species in the ionization reaction (Use ICE table if needed).
Use Ka and the equilibrium constant expression to find [H+].
Use [H+]eq and the initial acid concentration, [HA]0, to solve for percent ionization.
Find the percent ionization of a 2.5 M HNO2 solution. The Ka of nitrous acid is 4.6 x 10-4.
HNO2 ⇄ H+ + NO2-
𝑲 𝒂 = H + [ NO 𝟐 − ] [ HNO 𝟐 ]
= 𝒙 𝟐 𝟐.𝟓 −𝒙
(x is small)
𝟒.𝟔× 𝟏𝟎 −𝟒 = 𝒙 𝟐 𝟐.𝟓
Solve for x, which equals [H+]

67. Acid-base properties of salts
Salts are ionic compounds (cation + anion)
When dissolved in water, some salts are pH-neutral, some are acidic, and some are basic
IN GENERAL, anions tend to form basic or neutral solutions.
IN GENERAL, cations tend to form acidic or neutral solutions.

68. Anions as weak bases
Anions can be thought as the conjugate base of an acid (the acid lost its proton, H+)
This anion is the conjugate base of
this acid
Cl-
 HCl F-
 HF
NO3-
 HNO3
C2H3O2-
 HC2H3O2

69. Anions as weak bases
This anion is the conjugate base of
this acid
Cl-
 HCl F-
 HF
NO3-
 HNO3
C2H3O2-
 HC2H3O2
In general, the anion A- is the conjugate base of HA.
However, not every anion will act as a base…
Conjugate base of a weak acid, it’s a weak base
Conjugate base of a strong acid, it’s pH-neutral

70. Classify the anion as basic or pH-neutral:
NO3-
NO2-
C2H3O2-
Conj. Base of a strong acid ∴ pH-neutral
Conj. Base of a weak acid (HNO3) ∴ weak base
Conj. Base of a weak acid (HC2H3O2)
∴ weak base

71. Cations as weak acids Cations that are the counterions of strong bases
NaOH, KOH, and Ca(OH)2 are examples of strong bases, and form OH- in solution.
Cations Na+, K+, and Ca2+ do not contribute to acidity
THEREFORE, they are pH-neutral

72. Cations that are the conjugate acids of weak bases
A cation can be formed from any non-ionic weak base by adding a proton, H+.
Cations can be thought as the conjugate acid of an base (the base gained a proton, H+)
This cation is the conjugate acid of
this weak base
NH4+
 NH3
C2H5NH3+
 C2H5NH2
CH3NH3+
 CH3NH2

73. BH+(aq) + H2O(aq) ⇄ H3O+(aq) + B(aq)
This cation is the conjugate acid of
this weak base
NH4+
 NH3
C2H5NH3+
 C2H5NH2
CH3NH3+
 CH3NH2
Cations with the general formula BH+ will act as a weak acid according to the equation:
BH+(aq) + H2O(aq) ⇄ H3O+(aq) + B(aq)
Conjugate acid of a weak base, it’s a weak acid

74. Classify the cation as acidic or pH-neutral:
C5H5NH+
Ca2+
Cr3+
Conj. Acid of a weak base ∴ a weak acid
Conj. Acid of a strong base ∴ pH-neutral
Small, highly charged cation ∴ a weak acid

76. Strong Acid – Strong base
Reacts completely
Yields neutral solution (pH = 7)
Example: HCl + NaOH  H2O(l) + NaCl
Net ionic equation:
H+(aq) + OH-(aq)  H2O(l)

77. Strong Acid – Weak base Reacts completely
Yields a weak acid (pH < 7)
Weak acid is a product of the reaction (it’s the conj. acid of the strong base)
Example:
HCl + CH3COONa(aq)  CH3COOH(aq) + NaCl(aq)
Net Ionic Equation:
H+(aq) + CH3COO-(aq)  CH3COOH

78. CH3COOH(aq) + OH-(aq) CH3COO-(aq) + H2O(l)
Weak acid – strong base
Reacts completely
Yields a weak base (pH > 7)
Weak base is a product of the reaction (it’s the conj. base of the strong acid)
Example: CH3COOH(aq) + NaOH(aq)
Net Ionic Equation:
CH3COOH(aq) + OH-(aq) CH3COO-(aq) + H2O(l)

79. Weak acid – Weak base Does not react well
Yields a buffer solution (described later)
Both a weak acid & weak base exist together
Example: CH3COOH(aq) + CH3COO-(aq)  CH3COOH(aq) + CH3COO-(aq)
No net reaction

80. L acid
1 L
X mol acid (H+)
1 mol acid (H+)
1 mol base (OH-)
Y mol base (OH-)
L base
1 L
Y mol base (OH-)
1 mol base (OH-)
1 mol acid (H+)
X mol acid (H+)

81. Equilibrium lies to the side of the weaker acid and base

82. Acids Donate ONE proton at a time!
monoprotic--acids donating one H+ (ex. HC2H3O2)
diprotic--acids donating two H+'s (ex. H2C2O4)
polyprotic--acids donating many H+'s (ex. H3PO4)
polyprotic bases--accept more than one H+; anions with −2 and −3 charges (ex. PO43− ; HPO42−)
Amphiprotic or amphoteric --molecules or ions that can behave as EITHER acids or bases; water, anions of weak acids (look at the examples above—sometimes water was an acid, sometimes it acted as a base)

83. Autoionization of water (Video)
Add notes to page 7 using the video link above. Include examples given in the video.

84. Autoionization of water
H2O(l) ⇄ H+(aq) + OH-(aq)
2 H2O ⇄ OH- + H3O+
See video for notes

85. H+/OH- Equilibrium
See video for notes
Kw =

86. What is the ion-product constant? How does it relate to [H+] and [OH-]?
See video for notes

87. Acidic or basic or neutral?
It is important to recognize the meaning of Kw. In an aqueous solution at 25°C, no matter what it contains, [H+] x [OH-] = 1.0 x
There are three possible situations:
A neutral solution, where [H+] = [OH-].
An acidic solution, where [H+] > [OH-].
A basic solution, where [H+] < [OH-].

88. Other temperatures… (add to p. 7)
If at a temperature OTHER THAN 25°C, the Kw will be a different value.
Remember that the numerical value of Keq is dependent on temperature!

89. Example 1 𝑲 𝒘 = H + OH − =𝟏.𝟎× 𝟏𝟎 −𝟏𝟒 (since it’s at 25°C)
Calculate [H+] or [OH-] as required for each of the following solutions at 25°C, and state whether the solution is neutral, acidic, or basic.
1.0 x 10-5 M OH-
1.0 x 10-7 M OH-
𝟏.𝟎× 𝟏𝟎 −𝟏𝟒 = H + (𝟏.𝟎× 𝟏𝟎 −𝟓 )
[OH-] > [H+]
∴ basic
H + = 𝟏.𝟎× 𝟏𝟎 −𝟏𝟒 𝟏.𝟎× 𝟏𝟎 −𝟓 = 𝟏.𝟎× 𝟏𝟎 −𝟗
[OH-] = [H+]
∴ neutral
𝟏.𝟎× 𝟏𝟎 −𝟏𝟒 = H + (𝟏.𝟎× 𝟏𝟎 −𝟕 )
H + = 𝟏.𝟎× 𝟏𝟎 −𝟏𝟒 𝟏.𝟎× 𝟏𝟎 −𝟕 = 𝟏.𝟎× 𝟏𝟎 −𝟕

90. Example 1 𝑲 𝒘 = H + OH − =𝟏.𝟎× 𝟏𝟎 −𝟏𝟒 (since it’s at 25°C)
Calculate [H+] or [OH-] as required for each of the following solutions at 25°C, and state whether the solution is neutral, acidic, or basic.
10.0 M H+
𝟏.𝟎× 𝟏𝟎 −𝟏𝟒 =(𝟏𝟎.𝟎) OH −
[H+] > [OH-]
∴ acidic
OH − = 𝟏.𝟎× 𝟏𝟎 −𝟏𝟒 𝟏𝟎.𝟎 = 𝟏.𝟎× 𝟏𝟎 −𝟏𝟓

91. Example 2 At 60°C, the value of Kw is 1 x 10-13.
Using Le Chatelier’s principle, predict whether the reaction is exothermic or endothermic
2 H2O ⇄ H3O+ + OH-
At 25°C, Kw = 1 x 10-14
At a higher temperature, Kw = 1 x
Kw increases with temperature, energy is a reactant, and so the process must be endothermic.
(The greater the K, the more products you have.)

92. Example 2 At 60°C, the value of Kw is 1 x 10-13.
Calculate [H+] and [OH-] in a neutral solution at 60°C.
2 H2O ⇄ H3O+ + OH-
𝟏.𝟎× 𝟏𝟎 −𝟏𝟑 = H + O H −
𝟏.𝟎× 𝟏𝟎 −𝟏𝟑 = 𝒙 𝟐
𝟏.𝟎× 𝟏𝟎 −𝟏𝟑 = 𝒙 𝟐
𝟑× 𝟏𝟎 −𝟕 𝑴=𝒙= H + = O H −

93. At 10°C, the autoionization of water equals 2. 9 x 10-15
At 10°C, the autoionization of water equals 2.9 x Calculate [H+] and [OH-] as necessary under each of the following conditions:
Calculate [H+] if [OH-] = 9.3 x 10-4 M. Is the solution acidic or basic?
Calculate [H+] and [OH-] for a neutral solution.
Calculate [OH-] if [H+] = 6.7 x M. Is the solution acidic or basic?

94. pH Scale https://www.youtube.com/watch?v=pFK16GsU1e4
Add notes to page 9 using the video link above. Include examples given in the video.

95. pH Scale
See video for notes 7 14

96. Why does [H+] = [OH-] at pH 7?
See video for notes

97. pH and pOH
See video for notes pH
[H+] (M)
[OH-] (M)
pOH 7 6 3 12

98. pH and pOH What is the pattern between pH and [H+]?
What is the relationship between [H+] and [OH-]?
How is pH related to pOH?

99. Calculating pH https://www.youtube.com/watch?v=aw-Y1d2IIJc
Add notes to page 10 using the video link above. Include examples given in the video.

100. See video for notes pH = – log[H+] pOH = – log[OH–]
The pH of a solution is 3.75, what is the Hydrogen ion concentration, [H+]?

101. Examples
What is the pH of a solution with a hydrogen ion concentration of 2.0 x 10-11?
Calculate the pH of a 1.0 x 10-3 M OH- solution. (Hint: Use Kw to first find [H+]. Assume 25°C.)
What is the pOH of a solution with a hydroxide ion concentration of 1.8 x 10-4?
Calculate the pOH of a 1.0 M H+ solution. (Hint: Use Kw to first find [OH-]. Assume 25°C.)
The pH of human blood was measured to be 7.41 at 25°C. Calculate the pOH, [H+], and [OH-] for the sample.
10.70
11.00
3.74
14.00
pOH = 6.59, [H+] = 3.9 x 10-8 M, [OH-] = 2.6 x 10-7 M