1. HCl(aq)  H+(aq)+ Cl-(aq)
Definitions
Arrhenius - In aqueous solution…
Acids increase hydrogen ion concentration [H+]
HCl(aq)  H+(aq)+ Cl-(aq) H Cl O – +
HCl(aq) + H2O(l) H3O+(aq)+ Cl-(aq)
Courtesy Christy Johannesson

2. KOH(aq)  K+(aq)+ OH-(aq)
Definitions
Arrhenius - In aqueous solution…
Bases increase hydroxide ion concentration [OH-]
KOH(aq)  K+(aq)+ OH-(aq)
(aq)
(aq)
(aq)
Courtesy Christy Johannesson

3. A Brønsted-Lowry acid is a proton donor
A Brønsted-Lowry base is a proton acceptor
conjugate acid
conjugate base
base
acid

4. Conjugate Acids and Bases:
From the Latin word conjugare, meaning “to join together.”
Reactions between acids and bases always yield their conjugate bases and acids.

5. More examples of Conjugate Pairs

6. Strong and Weak Acids/Bases
The strength of an acid (or base) is determined by the amount of IONIZATION.
HNO3, HCl, H2SO4 and HClO4 are among the only known strong acids.

7. Strong and Weak Acids/Bases
Weak acids are much less than 100% ionized in water.
One of the best known is acetic acid = CH3CO2H

8. “Strong” = complete dissociation in solution
“Weak” = incomplete dissociation in solution

9. Strong and Weak Acids/Bases
Strong Base = 100% dissociated in water.
NaOH (aq) → Na+ (aq) + OH- (aq)
CaO
Other common strong bases include KOH and Ca(OH)2
CaO (lime) + H2O -->
Ca(OH)2 (slaked lime)

10. Strong and Weak Acids/Bases
Weak base = less than 100% ionized in water
One of the best known weak bases is ammonia
NH3 (aq) + H2O (l)  NH4+ (aq) + OH- (aq)

11. Extra (nontestable) information… Factors Affecting Acid Strength
The more polar the H-X bond and/or the weaker the H-X bond, the more acidic the compound.
Acidity increases from left to right across a row and from top to bottom down a group.

12. Yet another reason water is special
It’s called “amphoteric” because H2O can function as both an ACID and a BASE.
Equilibrium constant for water = Kw
Kw = [H+] [OH-] = x at 25 oC

13. Acid – Base Concentrations
pH = 3
pH = 11
10-1 H+
pH = 7
OH-
10-7
concentration (moles/L)‏ H+
OH-
OH- H+
10-14
[H+] > [OH-]
[H+] = [OH-]
[H+] < [OH-]
acidic
solution
neutral
solution
basic
solution
Timberlake, Chemistry 7th Edition, page 332

14. pH of Common Substance More basic More acidic 14 1 x 10-14 1 x 10-0 0
pH [H1+] [OH1-] pOH
14 1 x x
13 1 x x
12 1 x x
11 1 x x
10 1 x x
9 1 x x
8 1 x x
6 1 x x
5 1 x x
4 1 x x
3 1 x x
2 1 x x
1 1 x x
0 1 x x
NaOH, 0.1 M
Household bleach
Household ammonia
Lime water
Milk of magnesia
Borax
Baking soda
Egg white, seawater
Human blood, tears
Milk
Saliva
Rain
Black coffee
Banana
Tomatoes
Wine
Cola, vinegar
Lemon juice
Gastric juice
More basic
7 1 x x
More acidic

15. (also uses a logarithmic scale)
Richter Scale – Earthquakes 6 7 5 .
Area Radius Richter Scale Diameter pH [H+]
mm M M M M M M M 4 3 2 1

16. pH Calculations pH = -log[H+] pH [H+] = 10-pH pH + pOH = 14
Kw = [H+] [OH-] = 1 x10-14
The pH scale is a concise way of describing the H3O+ concentration and the acidity or basicity of a solution
• pH and H+ concentration are related as follows:
pH = –log10[H+] or [H+] = 10–pH
• pH of a neutral solution ([H3O+] = 1.00 x 10–7 M) is 7.00
• pH of an acidic solution is < 7, corresponding to [H3O+] > 1.00 x 10–7
• pH of a basic solution is > 7, corresponding to [H3O+] < 1.00 x 10–7
• The pH scale is logarithmic, so a pH difference of 1 between two solutions corresponds to a difference of a factor of 10 in their hydronium ion concentrations
There is an analogous pOH scale to describe the hydroxide ion concentration of a solution; pOH and [OH–] are related as follows:
pH = –log10[OH–] or [OH–] = 10–pOH
• A neutral solution has [OH–] = 1.00 x 10–7, so the pOH of a neutral solution is 7.00
• The sum of the pH and the pOH for a neutral solution at 25ºC is = 14.00
pKw = –log Kw = –log([H3O+] [OH–]) =
(–log[H3O+]) + (–log[OH–]) = pH + pOH
• At any temperature, pH + pOH = pKw, and at 25ºC, where Kw = 1.01 x 10–14, pH + pOH = 14.00; pH of any neutral solution is just half the value of pKw at that temperature
pOH
pOH = -log[OH-]
[OH-]
[OH-] = 10-pOH

17. (less acidic than vinegar,
What is the pH of a solution whose pOH is 11.09? Is this solution acidic or basic?
pH + pOH = 14
pH = 14
Solution is acidic 
(less acidic than vinegar,
more than orange juice)
pH = 14 – = 2.91

19. pH Calculations pH = -log[H+] pH [H+] = 10-pH pH + pOH = 14
Kw = [H+] [OH-] = 1 x10-14
The pH scale is a concise way of describing the H3O+ concentration and the acidity or basicity of a solution
• pH and H+ concentration are related as follows:
pH = –log10[H+] or [H+] = 10–pH
• pH of a neutral solution ([H3O+] = 1.00 x 10–7 M) is 7.00
• pH of an acidic solution is < 7, corresponding to [H3O+] > 1.00 x 10–7
• pH of a basic solution is > 7, corresponding to [H3O+] < 1.00 x 10–7
• The pH scale is logarithmic, so a pH difference of 1 between two solutions corresponds to a difference of a factor of 10 in their hydronium ion concentrations
There is an analogous pOH scale to describe the hydroxide ion concentration of a solution; pOH and [OH–] are related as follows:
pH = –log10[OH–] or [OH–] = 10–pOH
• A neutral solution has [OH–] = 1.00 x 10–7, so the pOH of a neutral solution is 7.00
• The sum of the pH and the pOH for a neutral solution at 25ºC is = 14.00
pKw = –log Kw = –log([H3O+] [OH–]) =
(–log[H3O+]) + (–log[OH–]) = pH + pOH
• At any temperature, pH + pOH = pKw, and at 25ºC, where Kw = 1.01 x 10–14, pH + pOH = 14.00; pH of any neutral solution is just half the value of pKw at that temperature
pOH
pOH = -log[OH-]
[OH-]
[OH-] = 10-pOH

20. pH = 0.620 Solution is acidic  pH = - log [H+] = -log [0.240M H+]
Calculate the pH of a M nitric acid solution. Is this solution acidic or basic?
pH = - log [H+] = -log [0.240M H+]
pH = 0.620
Solution is acidic 
(less acidic than battery acid,
more than stomach fluids)

21. 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
1.0 M
NaOH
(lye)‏
14.0
orange
3.5
potato
5.8
blood
7.4
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-]
Timberlake, Chemistry 7th Edition, page 335

22. pH Calculations pH = -log[H+] pH [H+] = 10-pH pH + pOH = 14
Kw = [H+] [OH-] = 1 x10-14
The pH scale is a concise way of describing the H3O+ concentration and the acidity or basicity of a solution
• pH and H+ concentration are related as follows:
pH = –log10[H+] or [H+] = 10–pH
• pH of a neutral solution ([H3O+] = 1.00 x 10–7 M) is 7.00
• pH of an acidic solution is < 7, corresponding to [H3O+] > 1.00 x 10–7
• pH of a basic solution is > 7, corresponding to [H3O+] < 1.00 x 10–7
• The pH scale is logarithmic, so a pH difference of 1 between two solutions corresponds to a difference of a factor of 10 in their hydronium ion concentrations
There is an analogous pOH scale to describe the hydroxide ion concentration of a solution; pOH and [OH–] are related as follows:
pH = –log10[OH–] or [OH–] = 10–pOH
• A neutral solution has [OH–] = 1.00 x 10–7, so the pOH of a neutral solution is 7.00
• The sum of the pH and the pOH for a neutral solution at 25ºC is = 14.00
pKw = –log Kw = –log([H3O+] [OH–]) =
(–log[H3O+]) + (–log[OH–]) = pH + pOH
• At any temperature, pH + pOH = pKw, and at 25ºC, where Kw = 1.01 x 10–14, pH + pOH = 14.00; pH of any neutral solution is just half the value of pKw at that temperature
pOH
pOH = -log[OH-]
[OH-]
[OH-] = 10-pOH

23. What is the hydroxide ion concentration of a solution with a hydronium ion concentration of 3.3 x 10-10M? Is this solution acidic or basic?
Kw = [H+] [OH-] = 1.0 x10-14
1.0 x10-14 = [3.3 x 10-10] [OH-]
[OH-] = 3.0 x 10-5 M
Solution is basic 
(less basic than ammonia,
more than sea water)

24. pH Calculations pH = -log[H+] pH [H+] = 10-pH pH + pOH = 14
Kw = [H+] [OH-] = 1 x10-14
The pH scale is a concise way of describing the H3O+ concentration and the acidity or basicity of a solution
• pH and H+ concentration are related as follows:
pH = –log10[H+] or [H+] = 10–pH
• pH of a neutral solution ([H3O+] = 1.00 x 10–7 M) is 7.00
• pH of an acidic solution is < 7, corresponding to [H3O+] > 1.00 x 10–7
• pH of a basic solution is > 7, corresponding to [H3O+] < 1.00 x 10–7
• The pH scale is logarithmic, so a pH difference of 1 between two solutions corresponds to a difference of a factor of 10 in their hydronium ion concentrations
There is an analogous pOH scale to describe the hydroxide ion concentration of a solution; pOH and [OH–] are related as follows:
pH = –log10[OH–] or [OH–] = 10–pOH
• A neutral solution has [OH–] = 1.00 x 10–7, so the pOH of a neutral solution is 7.00
• The sum of the pH and the pOH for a neutral solution at 25ºC is = 14.00
pKw = –log Kw = –log([H3O+] [OH–]) =
(–log[H3O+]) + (–log[OH–]) = pH + pOH
• At any temperature, pH + pOH = pKw, and at 25ºC, where Kw = 1.01 x 10–14, pH + pOH = 14.00; pH of any neutral solution is just half the value of pKw at that temperature
pOH
pOH = -log[OH-]
[OH-]
[OH-] = 10-pOH

25. Calculate the [H+] of a solution with a pH of 4.60
pH = - log [H+]
4.60 = -log [H+]
[H+] = 2.51 x 10-5 M
2nd
log
-4.6
You can check your answer by working backwards.
pH = - log [2.51x10-5 M]
pH = 4.60

26. pH Calculations pH = -log[H+] pH [H+] = 10-pH pH + pOH = 14
Kw = [H+] [OH-] = 1 x10-14
The pH scale is a concise way of describing the H3O+ concentration and the acidity or basicity of a solution
• pH and H+ concentration are related as follows:
pH = –log10[H+] or [H+] = 10–pH
• pH of a neutral solution ([H3O+] = 1.00 x 10–7 M) is 7.00
• pH of an acidic solution is < 7, corresponding to [H3O+] > 1.00 x 10–7
• pH of a basic solution is > 7, corresponding to [H3O+] < 1.00 x 10–7
• The pH scale is logarithmic, so a pH difference of 1 between two solutions corresponds to a difference of a factor of 10 in their hydronium ion concentrations
There is an analogous pOH scale to describe the hydroxide ion concentration of a solution; pOH and [OH–] are related as follows:
pH = –log10[OH–] or [OH–] = 10–pOH
• A neutral solution has [OH–] = 1.00 x 10–7, so the pOH of a neutral solution is 7.00
• The sum of the pH and the pOH for a neutral solution at 25ºC is = 14.00
pKw = –log Kw = –log([H3O+] [OH–]) =
(–log[H3O+]) + (–log[OH–]) = pH + pOH
• At any temperature, pH + pOH = pKw, and at 25ºC, where Kw = 1.01 x 10–14, pH + pOH = 14.00; pH of any neutral solution is just half the value of pKw at that temperature
pOH
pOH = -log[OH-]
[OH-]
[OH-] = 10-pOH

27. Why do chemists use titrations?
Quantitative analysis — used to determine the amounts or concentrations of substances present in a sample by using a combination of chemical reactions and stoichiometric calculations

28. Titration
standard solution
unknown solution
Definition
Analytical method in which a standard solution is used to determine the concentration of an unknown solution.
Quantitative analysis — used to determine the amounts or concentrations of substances present in a sample by using a combination of chemical reactions and stoichiometric calculations
Titration
– A method in which a measured volume of a solution of known concentration, called the titrant, is added to a measured volume of a solution containing a compound whose concentration is to be determined (the unknown)
– Reaction must be fast, complete, and specific (only the compound of interest should react with the titrant)
– Equivalence point — point at which exactly enough reactant has been added for the reaction to go to completion
Courtesy Christy Johannesson

29. Buret
stopcock
Erlenmeyer flask

30. Titration Vocabulary
Titrant
The substance added to the analyte in a titration (a standard solution)
Analyte
The substance being analyzed
Equivalence point
The point in a titration at which the quantity of titrant is exactly sufficient for stoichiometric reaction with the analyte.

31. Analyte Acid-Base Titration Titrant
If the concentration of the titrant is known, then the unknown concentration of the analyte can be determined.
Acid-Base Titration
Titrant
In acid-base titrations, a buret is used to deliver measured volumes of an acid or base solution of known titration (the titrant) to a flask that contains a solution of a base or an acid, respectively, of unknown concentration (the unknown).
If the concentration of the titrant is known, then the concentration of the unknown can be determined.
Plotting the pH changes that occur during an acid-base titration against the amount of acid or base added produces a titration curve; the shape of the curve provides important information about what is occurring in solution during the titration.
Before addition of any strong base, the initial [H3O+] equals the concentration of the strong acid.
Addition of strong base before the equivalence point, the point at which the number of moles of base (or acid) added equals the number of moles of acid (or base) originally present in the solution, decreases the [H3O+] because added base neutralizes some of the H3O+ present.
Addition of strong base at the equivalence point neutralizes all the acid initially present and pH = 7.00; the solution contains water and a salt derived from a strong base and a strong acid.
Addition of a strong base after the equivalence causes an excess of OH– and produces a rapid increase in pH.
A pH titration curve shows a sharp increase in pH in the region near the equivalence point and produces an S-shaped curve; the shape depends only on the concentration of the acid and base, not on their identity.
For the titration of a monoprotic strong acid with a monobasic strong base, the volume of base needed to reach the equivalence point can be calculated from the following relationship:
moles of base = moles of acid
(volume)b (molarity)b = (volume)a (molarity)a
VbMb = VaMa
Analyte

32. Buret Reading

33. Acidic, basic, or neutral??

34. The “perfect pink” for a titration with phenolphthalein
Phenolphthalein is clear in acidic solutions but pink in basic solutions!

35. point at which exactly enough reactant has been added for the solution to be neutralized and no more
Indicator - changes color to indicate pH change
pink pH
Example… phenolphthalein is colorless in acid and pink in basic solution
Endpoint = 7
colorless
Volume base added

36. Equivalence point (endpoint)
Point at which equal amounts of H+ and OH- have been added.
Determined by…
indicator color change
dramatic change in pH
Most common acids and bases are not intensely colored
– Rely on an indicator
Endpoint — point at which a color change is observed, which is close to the equivalence point in an acid-base titration
Courtesy Christy Johannesson

37. moles H+ = moles OH- Titration At endpoint...
Courtesy Christy Johannesson

38. MA = 0.66M HCl Titration Acid (H+) Base (OH-) M = ? M = 1.3M
25.5 mL of 1.3M KOH are required to neutralize mL of HCl. Find the molarity of HCl.
MA = 0.66M HCl
Acid (H+)
M = ?
V = 50.0 mL
Base (OH-)
M = 1.3M
V = 25.5 mL
A reaction in which an acid and a base react in stoichiometric amounts to produce water and a salt
Strengths of the acid and base determine whether the reaction goes to completion
1. Reactions that go to completion
a. Reaction of any strong acid with any strong base b. Reaction of a strong acid with a weak base
c. Reaction of weak acid with a weak base
2. Reaction that does not go to completion is a reaction of a weak acid or a weak base with water
Courtesy Christy Johannesson

39. Typical Example... pH Titration Data NaOH added (mL) pH
Titration of an Acid With an Base
14.0
phenolphthalein - pink
12.0
10.0
8.0 pH
equivalence point
6.0
4.0
phenolphthalein - colorless
Solution
of NaOH
Solution
of NaOH
2.0
Na+
OH-
0.0
0.0
10.0
20.0
30.0
40.0
Volume of M NaOH added
(mL) H+
Cl-
Solution
of HCl
25 mL