1. Acids and Bases ACIDS and BASES
“In nature, acids can be found in fruits: citric acid is responsible for the sharp taste of lemons. Vinegar contains acetic acid, and tannic acid from tree bark is used to tan leather. The stronger mineral acids have been prepared since the Middle Ages. One of these, aqua fortis (nitric acid), was used by assayers to separate gold from silver. Car batteries contain sulfuric acid, also strong and corrosive. A base is the opposite of an acid. Bases often feel slippery; bicarbonate of soda and soap are bases, and so is lye, a substance that can burn skin. Bases that dissolve in water are called alkalis. In water, acids produce hydroxide ions. When an acid and a base react together, the hydrogen and hydroxide ions combine and neutralize each other, forming water together and a salt. The strength of acids and bases can be measured on a pH scale.”
Eyewitness Science “Chemistry” , Dr. Ann Newmark, DK Publishing, Inc., 1993, pg 42

2. Concentration vs. Strength
Concentrated vs. Dilute
Strong vs. Weak

3. A- H+ A- H+ A- H+ A- H+ A- H+
Acids: Concentration vs. Strength
CONCENTRATED
DILUTE
H+ A- H+ A- H+ A- H+ A- HA
A- H+ A- H+ A- H+ A- H+ A -
H+ A- HA H+ A- H+ A- H+ A-
A- H+ A- H+ A- H+ A- H+ A- H+
H+ A - H + A - H + A - HA H + A -
A- H+ A- H+ A- H+ A- H+ A–
H+ A- H+ A- H+ A- H+ A- H+
A- H+ A- H+ A- H+ A- H+ A-
HA A- H+ A- H+ A- H+ A- H+
H+ A H+ A HA
A H+ A - H A –
H+ A H+ A H+
A HA H A -
H+ A H+ A H+
STRONG ACIDS
Dissociate nearly 100%
HA H A-
WEAK ACIDS
Dissociate very little
HA H A-
STRONG
HA HA H+ A- HA HA HA
HA HA HA HA HA H+ A-
H+ A- HA HA HA HA HA
HA HA H+ A- HA HA HA
HA HA HA H+ A- HA HA
H+ A- HA HA HA HA HA
HA HA HA H+ A- HA HA
H+ A- HA HA HA HA HA
HA HA H+ A- HA HA HA
HA HA HA HA
HA HA HA
H+ A HA HA
HA HA H + A –
HA H + A – HA HA
Strong acids react essentially completely with water to give H+ and the corresponding anion.
Strong bases dissociate essentially completely in water to give OH– and the corresponding cation.
Both strong acids and strong bases are strong electrolytes.
Only a fraction of the molecules of weak acids and weak bases react with water to produce ions and are, therefore, weak electrolytes.
No correlation between solubility in water and whether a substance is a strong or a weak electrolyte.
WEAK

4. Strength DOES NOT mean concentration
Strong vs. Weak Acid
Strength DOES NOT mean concentration
A salt can dissolve in water to produce a neutral, basic, or acidic solution, depending on whether it contains the conjugate base of a weak acid as the anion (A–) or the conjugate acid of a weak base as the cation (BH+), or both.
• Salts that contain small, highly charged metal ions produce acidic solutions in water.
• The most important parameter for predicting the effect of a metal ion on the acidity of coordinated water molecules is the charge-to-radius ratio of the metal ion.
• The reaction of a salt with water to produce an acidic or basic solution is called a hydrolysis reaction, which is just an acid-base reaction in which the acid is a cation or the base is an anion.
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 508

5. pH Scale Søren Sorensen (1868 - 1939)
The pH scale was invented by the Danish chemist Søren Sorensen for a brewery to measure the acidity of beer.
Søren Sorensen
( )

6. pH Scale pH is a way to measure the strength of an acid or base. 7
[H+] pH
1 M NaOH
Ammonia
(household
cleaner)
Blood
Pure water
Milk
Vinegar
Lemon juice
Stomach acid
1 M HCl
Acidic Neutral Basic
Acid
Base 7 14
pH Scale
pH is a way to measure the strength of an acid or base.
“S.P.L. Sorensen ( ) introduced the pH scale to measure the concentration of hydrogen ions in solution. The more hydrogen ions, the stronger the acid. The amount of hydrogen ions in solution can affect the color of certain dyes found in nature. These dyes can be used as indicators to test for acids and alkalis. An indicator such as litmus (obtained from lichen) is red in acid. If base is slowly added, the litmus will turn blue when the acid has been neutralized, at about 6-7 on the pH scale. Other indicators will change color at different pH’s. A combination of indicators is used to make a universal indicator.”
- Eyewitness Science “Chemistry” , Dr. Ann Newmark, DK Publishing, Inc., 1993, pg 42
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 515

7. 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

8. pH of Common Substance 14 1 x 10-14 1 x 10-0 0 13 1 x 10-13 1 x 10-1 1
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

9. Acid – Base Concentrations
10-1
pH = 3
pH = 11
H3O+
OH-
pH = 7
10-7
concentration (moles/L)
H3O+
OH-
OH-
H3O+
10-14
[H3O+] > [OH-]
[H3O+] = [OH-]
[H3O+] < [OH-]
acidic
solution
neutral
solution
basic
solution
Timberlake, Chemistry 7th Edition, page 332

10. pH
pH = -log [H1+]
Kelter, Carr, Scott, Chemistry A World of Choices 1999, page 285

11. pH Calculations pH pOH [H3O+] [OH-] pH = -log[H3O+] [H3O+] = 10-pH
[H3O+] [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
[OH-]
pOH = -log[OH-]
[OH-] = 10-pOH

12. pH = - log [H+] Given: pH = 4.6 determine the [hydronium ion]
choose proper equation
4.6 = - log [H+]
substitute pH value in equation
= log [H+]
multiply both sides by -1
2nd
log
= log [H+]
take antilog of both sides
[H+] = 2.51x10-5 M
10x
antilog
Recall, [H+] = [H3O+]
You can check your answer by working backwards.
pH = - log [H+]
pH = - log [2.51x10-5 M]
pH = 4.6

13. pH Indicators
Substances that change color depending on the pH of the solution..

14. Litmus Paper
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

15. pH Paper pH pH

16. Paper or plastic strips that contain combinations of indicators estimate the pH of a solution by simply dipping a piece of pH paper into it and comparing the resulting color with standards printed on the container
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

17. Desired Features of Sensors
pH paper pH pH
Detection limit
Low deflection
High sensitivity
High selectivity
Wide dynamic range
Simple to use
Cost-effective
Most acid-base titrations are not monitored by recording the pH as a function of the amount of the strong acid or base solution used as a titrant
Instead, an acid-base indicator is used, and they are compounds that change color at a particular pH and if carefully selected, undergo a dramatic color change at the pH corresponding to the equivalence point of the titration
Acid-base indicators are typically weak acids or bases whose changes in color correspond to deprotonation or protonation of the indicator itself
The chemistry of indicators are described by the general equation Hn(aq) ⇋ H+ (aq) + n–(aq), where the protonated form is designated by Hn and the conjugate base by n–
The ionization constant for the deprotonation of indicator Hn is Kin = [H+] [n–] / [Hn]
The value of pKin determines the pH at which the indicator changes color A good indicator must have the following properties:
1. Color change must be easily detected
2. Color change must be rapid
3. Indicator molecule must not react with the substance being titrated
4. The indicator should have a pKin that is within one pH unit of the expected pH at the equivalence point of the titration
• Synthetic indicators have been developed that meet the above criteria and cover the entire pH range
• An indicator does not change color abruptly at a particular pH but undergoes a pH titration like any other acid or base

18. Range and Color Changes of Some Common Acid-Base 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
From F. Brescia et al., Chemistry: A Modern Introduction, W. B. Saunders Co., 1978.
Adapted from R. Bates, Determination of pH, Theory and Practice, John Wiley & Sons, Inc., New York, 1964.
Choosing the correct indicator for an acid-base titration
1. For titrations of strong acids and strong bases (and vice versa), any indicator with a pKin between 4 and 10 will do
2. For the titration of a weak acid, the pH at the equivalence point is greater than 7, and an indicator such as phenolphthalein or thymol blue, with pKin > 7, should be used
3. For the titration of a weak base, where the pH at the equivalence point is less than 7, an indicator such as methyl red or bromcresol blue, with pKin < 7, should be used
Phenolphthalein colorless red colorless beyond 13.0
Bromthymol blue indicator would be used in titrating a strong acid with a strong base.
Phenolpthalein 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.

19. pH
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

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

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

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

23. pH Indicator 1 2 3 4 5 6 7 8 9 10 11 12 Phenolphthalein Methyl Red
Colorless
Pink
Red
Methyl Red
Red
Orange
Yellow
Orange IV
Orange
Peach
Yellow
phenolphthalein
methyl red
methyl orange

24. Indicator Colors in Titration

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

26. 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 *
Source: Volume 62, Number 4, April 1985 pg 285 (Not sure of magazine title)
*YELLOW at pH 12 and above

27. Red Cabbage Indicator
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

28. Phenolphthalein Indicator
Colorless = Acidic pH
Pink = Basic pH

29. (Colorless acid form, HIn) (Pink base form, In-)OH HO C O O-
(Colorless acid form, HIn)
(Pink base form, In-)