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. Terms Acids Examples: Bases Compounds that form H+ in water.
Formulas usually begin with ‘H’.
Examples:
HCl – hydrochloric acid
HNO3 – nitric acid
H2SO4 – sulfuric acid
Bases
Compounds that form OH- in water
Formulas usually end in ‘hydroxide’

3. Properties ACIDS BASES electrolytes electrolytes sour taste
bitter taste
turn litmus red
turn litmus blue
react with metals to form H2 gas
slippery feel
vinegar, milk, soda, apples, citrus fruits
ammonia, lye, antacid, baking soda
ChemASAP

4. Acid + Base  Salt + Water
Orange juice + milk  bad taste
Evergreen shrub + concrete  dead bush
HCl + NaOH  NaCl + HOH
salt water

5. Common Acids and Bases 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
Phosphoric acid, H3PO4, is another acid commonly found in the laboratory. It is on the borderline between a strong and weak acid.
Weak Base (weak electrolyte)
NH4OH ammonia
Weak Acids (weak electrolytes)
CH3COOH acetic acid
H2CO3 carbonic
Weak Base (weak electrolyte)
NH ammonia
NH3 + H2O  NH4OH
Kotz, Purcell, Chemistry & Chemical Reactivity 1991, page 145

6. Acid Nomenclature
A binary acid – is an acid that contains only two different elements: hydrogen and one of the more electronegative elements.
Rules of naming a binary acid:
The name of the binary acid begins with the prefix hydro-
The root of the name of the second element follows this prefix.
The name then ends with the suffix –ic
A oxyacid – is an acid that is a compound of hydrogen, oxygen, and a third element, usually a nonmetal.

7. Acid Nomenclature

8. Acid Nomenclature HBr H2CO3  hydrobromic acid H2SO3  carbonic acid
2 elements, -ide
 hydrobromic acid
3 elements, -ate
 carbonic acid
3 elements, -ite
 sulfurous acid

9. Acid Nomenclature hydrofluoric acid sulfuric acid nitrous acid  H+ F-
2 elements
 H+ F-
 HF
3 elements, -ic
 H+ SO42-
 H2SO4
3 elements, -ous
 H+ NO2-
 HNO2

10. Common Industrial Acids
Sulfuric Acid
Petroleum refining
Metallurgy
Manufacture of manure
Essential to the processes of: production of metals, paper, paint, dyes, detergents and in automobile batteries.
Effective dehydration agent.
Nitric Acid
Is a volatile, unstable liquid rarely used in industry. Dissolving in water provides stability.
Used in the making of explosives, rubber, plastics, dyes and pharmaceuticals.
Stains proteins yellow.

11. Common Industrial Acids
Phosphoric Acid
Used directly for manufacturing of fertilizers and animal feed, detergents and ceramics.
Flavoring of beverages.
Cleaning agent for dairy equipment.
Hydrochloric Acid
Aids in digestion.
Used for “pickling” iron and steel.
General cleaning agent.
Used for food processing, activation of oil wells, recovery of Mg from sea water, and the production of other chemicals.

12. Common Industrial Acids
Acetic Acid
Used for synthesizing chemicals used in the manufacture of plastics.
Used as a fungicide.
Food supplements

13. pH of Common Products

14. : measures acidity/basicity
pH scale
: measures acidity/basicity
ACID
BASE
10x
100x
10x
10x
NEUTRAL
Each step on pH scale represents a factor of 10.
pH 5 vs. pH (10X more acidic)
pH 3 vs. pH (100X different)
pH 8 vs. pH (100,000X different)

15. : measures acidity/basicity
pH scale
: measures acidity/basicity
Soren Sorensen
( )
ACID
BASE
10x
100x
10x
10x
NEUTRAL
Each step on pH scale represents a factor of 10.
pH 5 vs. pH (10X more acidic)
pH 3 vs. pH (100X different)
pH 8 vs. pH (100,000X different)

16. Kw = [H3O+][OH-] = 1.0  10-14 H2O + H2O H3O+ + OH-
Ionization of Water
H2O + H2O H3O+ + OH-
Kw = [H3O+][OH-] = 1.0  10-14

17. Ionization of Water
Find the hydroxide ion concentration of 3.0  10-2 M HCl.
[H3O+][OH-] = 1.0  10-14
[3.0  10-2][OH-] = 1.0  10-14
[OH-] = 3.3  M
Acidic or basic?
Acidic

18. The pH Scale
pH – of a sol’n is the negative of the common logarithm of the hydronium ion concentration [H3O+].
pOH - of a sol’n is the negative of the common logarithm of the hydroxide ion concentration [OH-].

19. pouvoir hydrogène (Fr.)
pH Scale 14 7
INCREASING
ACIDITY
INCREASING
BASICITY
NEUTRAL
pH = -log[H3O+]
pouvoir hydrogène (Fr.)
“hydrogen power”

20. pH = -log[H3O+] pOH = -log[OH-] pH + pOH = 14
pH Scale
pH = -log[H3O+]
pOH = -log[OH-]
pH + pOH = 14

21. pH Scale What is the pH of 0.050 M HNO3? pH = -log[H3O+]
Acidic or basic?
Acidic

22. pH Scale
What is the molarity of HBr in a solution that has a pOH of 9.6?
pH + pOH = 14
pH = 14
pH = 4.4
pH = -log[H3O+]
[H3O+] = 10-pH
[H3O+] =
[H3O+] = 4.0  10-5 M HBr
Acidic

23. The pH Scale Basic pH and pOH calculations: pH = -log[H3O+]
pOH = -log[OH-]
pH = 14 – pOH
pOH = 14 – pH
Calculating pH from [H3O+]
Determine concentration of [H3O+] using Kw formula.
Determine pH using logarithm formula.
Calculating pOH from [OH-]
Determine concentration of [OH-] using Kw formula.
Determine pOH using logarithm formula.

24. Example pH (pOH) calculations
If you had a sol’n with the concentration of M [H3O+], what is the sol’ns pH? What is the sol’ns pOH?
1. Given: [H3O+] = M
Unknown: pH = ??
pH = -log [H3O+]
= -log[0.025M]
= 1.6
pOH = 14 – pH =
= 12.4

25. Example pH (pOH) calculations
B. What is the pH of M HNO3? Is it acidic or basic?
pH = -log[H3O+]
pH = -log[0.050] = 1.3
It is acidic because 1.3  7.

26. Example pH (pOH) calculations
C. What is the pH of a M HCl sol’n?
Given: Identity
Concentration: [0.100 M]
Unknown: pH = ????
2. [H3O+] = M
pH = -log[0.100 M] = 1.00

27. Example pH (pOH) calculations
D. What is the pH of a 1.5 M sol’n of KOH?
1. Given: Identity
Concentration: [1.5 M]
Unknown: pH = ????
2. [OH-] = 1.5 M
[H3O+] = 1.0 x M2/[OH-]
= 1.0 x M2/[1.5 M]
= 6.7 x 10-15M
pH = -log[H3O+]
= -log[6.7 x 10-15M]
= 14.17
Check : Answer does indicate that KOH forms a sol’n
pH  7, which is basic.

28. Example pH (pOH) calculations
E. Determine the hydronium concentration and the hydroxide concentration of an aqueous sol’n that has a pH of 5.6.
1. Given: pH = 5.6
Unknown: [H3O+] = ????
[OH-] = ????
2. [H3O+] = 10-pH
= = 2.5 x 10-6 M
[OH-] = 1.0 x M2/[H3O+]
= 1.0 x M2/[2.6 x 10-6] = 3.8 x 10-9 M
Check: The pH is acidic, therefore it should be found that the [H3O+]  [OH-]. Yes it is.

29. Example pH (pOH) calculations
F. What is the molarity of HBr in a solution that has a pOH of 9.6?
1. Given: pOH = 9.6
Unknown: pH = ????
[H3O+] = ????
2. pH + pOH = 14
pH = 14 – 9.6 = 4.4
[H3O+] = 10-pH = = 4.0 x 10-5 M

30. The Three Acid/Base Theories
Arrhenius
Bronsted-Lowry
Lewis

31. A Comparison b/w acid-base theories
Type
Acid
Base
Arrhenius
H+ or H3O+ producer
OH- producer
Bronsted-Lowry
Proton donor (H+)
Proton acceptor (H+)
Lewis
Electron-pair acceptor
Electron-pair donor

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

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

34. HCl + H2O  Cl– + H3O+ B. 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

35. B. Definitions
H2O + HNO3  H3O+ + NO3– B A CA CB

36. Amphoteric - can be an acid or a base.
B. Definitions
NH3 + H2O  NH4+ + OH- B A CA CB
Amphoteric - can be an acid or a base.

37. Polyprotic - an acid with more than one H+
B. Definitions
Give the conjugate base for each of the following: HF
H3PO4
H3O+
F -
H2PO4-
H2O
Polyprotic - an acid with more than one H+

38. Br - HSO4- CO32- HBr H2SO4 HCO3- B. Definitions
Give the conjugate acid for each of the following:
Br -
HSO4-
CO32-
HBr
H2SO4
HCO3-

39. 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
Bromothymol 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.
Phenolphthalein colorless red colorless beyond 13.0

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

41. 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

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

43. Acid-Base Neutralization1+ 1- + +
Hydronium ion
Hydroxide ion
H3O+
OH-
Water
H2O
Water
H2O
Dorin, Demmin, Gabel, Chemistry The Study of Matter 3rd Edition, page 584

44. Acid-Base Neutralization1+ 1- + +
H3O+
OH-
H2O
H2O
Hydronium ion
Hydroxide ion
Water
Water
Dorin, Demmin, Gabel, Chemistry The Study of Matter 3rd Edition, page 584

45. Titration Acids and bases sometimes react in a 1:1 mole ratio.
HCL(aq) + NaOH(aq)→ NaCl(aq) + H2O(l)
When sulfuric acid and sodium hydroxide react…the ratio is 1:2.
H2SO4(aq) + 2NaOH(aq) → Na2SO4(aq) + 2H2O(l)
This acid - base mole ratio from the balanced equation is a conversion factor used when neutralizing an acid or a base

46. Titration – using mole ratios…
How many moles of sulfuric acid would you require to neutralize 0.50 mol of sodium hydroxide?
STEP 1: Write the balanced equation for the rxn to find acid:base mol ratio
H2SO4(aq) + 2NaOH(aq) → Na2SO4(aq) + 2H2O(l)
STEP 2: Use the mole ratio to calculate the necessary number of moles of H2SO4
0.50 mol NaOH x 1 mol H2SO4 = 0.25 mol H2SO4
2 mol NaOH

47. Titration – Your Turn 
How many moles of sodium hydroxide are required to neutralize 0.20 mol of nitric acid?
STEP 1: NaOH + HNO3 → NaNO3 + H2O
STEP 2: 0.20 mol HNO3 x 1 mol NaOH
1 mol HNO3
ANSWER: 0.20 mol NaOH needed
How many moles of potassium hydroxide are needed to neutralize 1.56 mol of phosphoric acid?

48. Acid-Base Titration

49. 1: 30 mL of 0. 10M NaOH neutralized 25. 0mL of hydrochloric acid
1: 30 mL of 0.10M NaOH neutralized 25.0mL of hydrochloric acid. Determine the concentration of the acid
Write the balanced chemical equation for the reaction
Extract the relevant information from the question: NaOH      V = 30mL , M = 0.10M       HCl    V = 25.0mL, M = ?
Calculate moles NaOH       n(NaOH) = M x V =
From the balanced chemical equation find the mole ratio  NaOH:HCl      
Find moles HCl
Calculate concentration of HCl using the formula: M=n÷V

50. 2. What volume of 0. 75M H2SO4 is required to neutralize 25. 0 mL of 0
2. What volume of 0.75M H2SO4 is required to neutralize 25.0 mL of 0.427M KOH?
Write the balanced chemical equation for the reaction
Extract the relevant information from the question: KOH      V = 25.0mL , M = 0.427M       HCl    V = ?mL, M= 0.75M
Calculate moles KOH       n(KOH) = M x V =
From the balanced chemical equation find the mole ratio  KOH: H2SO4      
Find moles H2SO4
Calculate volume of H2SO4 using the formula: V=n÷M

51. 3. 50. 0 mL of an unknown solution of Ca(OH)2 are titrated with 0
mL of an unknown solution of Ca(OH)2 are titrated with 0.15M HCl. Find the molarity of the Ca(OH)2 solution if 83 mL of acid are required to reach the equivalence point.
Write the balanced chemical equation for the reaction
Extract the relevant information from the question: Ca(OH)2   V = 50.0mL , M = ?M       HCl    V = 83 mL, M= 0.15M
Calculate moles HCl       n(HCl) = M x V =
From the balanced chemical equation find the mole ratio  Ca(OH)2 : HCl      
Find moles Ca(OH)2
Calculate molarity of Ca(OH)2 using the formula: M=n÷V

52. Titration Curve

53. Titration – Virtual Lab