1. ACIDS AND BASES

2. Acids and Bases
Two important classification of compounds - Acids and Bases
Properties of ACIDS
Taste Sour/Tart
Stings and burns the skin
Reacts with bases
Turns blue litmus paper red
Reacts with metals to form H2 gas
Neutralizes Bases
Donates H+
Conduct electricity.
Can be strong or weak electrolytes in aqueous solution
Properties of BASES
Taste Bitter
Feels slippery on skin
Reacts with acids
Turns red litmus blue
Doesn’t react with metals
Neutralizes Acids
Accepts H+
Chapters 19 and 9.4

3. Acids Affect Indicators
Blue litmus paper turns red in contact with an acid.
Chapters 19 and 9.4

4. Acids React with Active Metals
Acids react with active metals to form salts and hydrogen gas:
HCl(aq) + Mg(s) → MgCl2(aq) + H2(g)
Chapters 19 and 9.4

5. Sulfuric Acid – H2SO4
Highest volume production of any chemical in the U.S.
Used in the production of paper
Used in production of fertilizers
Used in petroleum refining
Chapters 19 and 9.4

6. Nitric Acid – HNO3 Used in the production of fertilizers
Used in the production of explosives
Nitric acid is a volatile acid – its reactive components evaporate easily
Stains proteins (including skin!)
Chapters 19 and 9.4

7. Hydrochloric Acid - HCl
Used in the “pickling” of steel
Used to purify magnesium from sea water
Part of gastric juice, it aids in the digestion of proteins
Sold commercially as “Muriatic acid”
Chapters 19 and 9.4

8. Phosphoric Acid – H3PO4 A flavoring agent in sodas
Used in the manufacture of detergents
Used in the manufacture of fertilizers
Not a common laboratory reagent
Chapters 19 and 9.4

9. Acetic Acid – HC2H3O2 Used in the manufacture of plastics
Used in making pharmaceuticals
Acetic acid is the acid present in household vinegar
Chapters 19 and 9.4

10. Bases Affect Indicators
Red litmus paper turns blue in contact with a base.
Phenolphthalein turns purple in a base.
Chapters 19 and 9.4

11. Sodium hydroxide (lye), NaOH
Examples of Bases
Sodium hydroxide (lye), NaOH
Potassium hydroxide, KOH
Magnesium hydroxide, Mg(OH)2
Calcium hydroxide (lime), Ca(OH)2
Chapters 19 and 9.4

12. Bases Neutralize Acids
Milk of Magnesia contains magnesium hydroxide, Mg(OH)2, which neutralizes stomach acid, HCl.
2 HCl + Mg(OH)2
MgCl2 + 2 H2O
Chapters 19 and 9.4

13. Acid-Base Theories Arrhenius (1883) NaOH HCl H2SO4 CO32- NH3 H2O
Brønsted–Lowry (1923)
BF3
AlI3
Lewis ( )

14. Arrhenius Definition of A & B
An acid is a substance that dissociates in water to produce hydrogen ions (H+)
Example: Hydrochloric Acid (HCl)
A base is a substance that dissociates in water to produce hydroxide ions (OH-)
Example: Sodium Hydroxide (NaOH)

15. What happens when placed in water???
Chapters 19 and 9.4

16. Svante Arrhenius ( )
Chapters 19 and 9.4

17. Brønsted-Lowry Definition
An acid is any substance that can donate H+ ions.
A base is any substance that can accept H+ ions.
Expansion on Arrhenius definition of A&B
Defines A&B as not necessarily in water solution
Does not contain OH- ; covers bases such as ammonia NH3
Chapters 19 and 9.4

18. Johannes Bronsted Thomas Lowry (1879-1947) (1874-1936)
Chapters 19 and 9.4

19. Brønsted-Lowry Definition
H+ is really only just a proton (no electrons or neutrons), so definition is often in terms of protons
Brønsted-Lowry Base is a proton acceptor
Brønsted-Lowry Acid is a proton donor
Monoprotic Acids can only donate 1 H+: HCl
Diprotic Acids can donate 2 H+: H2SO4
Triprotic Acids can donate 3 H+: H3PO4
Polyprotic acids = diprotic and triprotic acids

20. Chapters 19 and 9.4

21. Chapters 19 and 9.4

22. Conjugate Acid-Base Pairs
Acid-base reactions with water proceed in both directions:
Example: NH3 (g) + H2O (l) NH4+ (aq) + OH- (aq)
Acid loses H+ = conjugate base
Example: H2O (acid) loses its H+, turning it into OH- (conjugate base)
Base gains H+ = conjugate acid
Example: NH3 (base) gains an H+, turning it into NH4+ (conjugate acid)
Chapters 19 and 9.4

23. Acids and bases come in pairs
A “conjugate base” is the remainder of the original acid, after it donates it’s hydrogen ion
A “conjugate acid” is the particle formed when the original base gains a hydrogen ion
Indicators are weak acids or bases that have a different color from their original acid and base
Chapters 19 and 9.4

24. The Hydronium Ion Water can pick up a H+ ion to form a hydronium ion:
H+ + H2O ® H3O H3O+ = hydronium ion
With acids, water is a Brønsted-Lowry base (accepts protons
Example: HCl (g) + H2O(l) ® H3O+ (aq) + Cl- (aq)
With bases, water is a Brønsted-Lowry acid (donates protons)
Example: NH3 (g) + H2O (l) ® NH4+ (aq) + OH- (aq)
Compound that can act as either a proton donor or acceptor = amphoteric
Chapters 19 and 9.4

25. Lewis Acids and Bases
Gilbert Lewis ( )
Gilbert Lewis focused on the donation or acceptance of a pair of electrons during a reaction
Lewis Acid - electron pair acceptor
Lewis Base - electron pair donor
Most general of all 3 definitions; acids don’t even need hydrogen!

26. Lewis Acids and Bases
Lewis acid = accepts a pair of electrons during a reaction
Lewis base = donates a pair of electrons during a reaction
Covers acids and bases not covered by Brønsted-Lowry definition
Table 19.4, Pg. 592, Text
Type
Acid
Base
Arrhenius
H+ producer
OH- producer
Brønsted-Lowry
H+ (proton) donor
H+ acceptor
Lewis
Electron-pair acceptor
Electron-pair donor
Chapters 19 and 9.4

27. The pH Scale
pH is based on the concentration of the hydronium ion in a solution
Concentrations range from 100 M (strong) to M (weak) of [H3O+]
pH ranges from 0 to 14
If [H30+] concentration = M, pH = 2
If [H30+] concentration = M, pH = 10
Water is 10-7 M, pH = 7
pH of 0-6 = Acidic
pH of 7 = Neutral
pH of 8-14 = Basic
Chapters 19 and 9.4

28. pH
…..defined as the negative base-10 logarithm of the hydronium ion concentration.
pH = −log [H3O+]
For pure H2O: [1.0  10−7 ]
= 7.0
Problem: Calculate pH when [H3O+] = 2.3 x 10-3 M
10x
log
log * 10^- 3
= 2.64
base 10 log
Problem: Calculate [H3O+] when pH = 2.3 ?
antilog
10x
log
10^
= 5.0 * 10-3

29. pH and pOH Calculations

30. Chapters 19 and 9.4

31. pH
These are the pH values for several common substances.
pH + pOH = 14

32. pH and pOH equilibrium in pure Water
[H3O+] [OH−] = Kw
[1.0  10−7 ] [1.0  10−7 ] = 1.0  10−14
−log [H3O+] + −log [OH−] = −log Kw
pH + pOH = pKw
Because in pure water [H3O+] = [OH−],
Kw = [1.0  10−7 ] [1.0  10−7 ] = 1.0  10−14 =

33. pH and pOH equilibrium in Water to which Acids & Bases are Added
Add base
OH-
H2O(l) + H2O(l)
H3O+(aq) + OH−(aq)
[H3O+] [OH−]
H2O
Kw = [1.0  10−7 ] [1.0  10−7 ] = 1.0  10−14
Kw = [1.0  10−8 ] [1.0  10−6 ] = 1.0  10−14
pH + pOH = pKw
= 14

34. Measuring pH Why measure pH? Solutions we use –
swimming pools
soil conditions for plants
medical diagnosis
soaps and shampoos, etc.
Sometimes we can use indicators, other times we might need a pH meter
Chapters 19 and 9.4

35. Acid-Base Strength
An acid or a base is considered strong if they completely dissociate into ions (H+ and OH-) in water
Strong Acids
HCl and H2SO4
Strong Bases
Hydroxides, e.g. NaOH
Conjugate acid-base pairs have an inverse relationship (works for both acids and bases)
The stronger the acid, the weaker the conjugate base
The weaker the acid, the stronger the conjugate base
Chapters 19 and 9.4

36. Strong Acids and Bases to know

37. Measuring pH with wide-range paper
1. Moisten indicator strip with a few drops of solution, by using a stirring rod.
2.Compare the color to the chart on the vial – read the pH value.
Chapters 19 and 9.4

38. Acid-Base Indicators
Although useful, there are limitations to indicators:
usually given for a certain temperature (25 oC), thus may change at different temperatures
what if the solution already has color, like paint?
the ability of the human eye to distinguish colors is limited
Chapters 19 and 9.4

39. Some of the many pH Indicators and their ranges
Chapters 19 and 9.4

40. Red Cabbage Juice as an indicator
Red cabbage juice mixed with baking soda (left) and with vinegar (right). On the top, a drop of unmixed juice.
Chapters 19 and 9.4

41. Acid-Base Indicators A pH meter may give more definitive results
some are large, others portable
works by measuring the voltage between two electrodes; typically accurate to within 0.01 pH unit of the true pH
needs to be calibrated
Chapters 19 and 9.4

42. Acids Neutralize Bases
HCl + NaOH → NaCl + H2O
-Neutralization reactions ALWAYS produce a salt and water.
-Of course, it takes the right proportion of acid and base to produce a neutral salt
Chapters 19 and 9.4

43. Acid-Base Properties of Salts
Neutralization reaction = reaction of an acid and a base
Acid + base react – form a salt and water
Type of salt depends on reactants
Acid
Base
Salt
Strong
Neutral
Weak
Acidic
Basic
Neutral, basic, or acidic

44. Chapters 19 and 9.4

53. Chapters 19 and 9.4

55. Buffers
Buffers are solutions in which the pH remains relatively constant, even when small amounts of acid or base are added
made from a pair of chemicals: a weak acid and one of it’s salts; or a weak base and one of it’s salts
Chapters 19 and 9.4

56. Chapters 19 and 9.4

57. Acid rain
Chapters 19 and 9.4

58. Causes of emissions
Chapters 19 and 9.4

59. pH readings nationwide
Chapters 19 and 9.4

60. Acid rain effects limestone and Marble
Chapters 19 and 9.4

61. Effects of Acid Rain on Marble (calcium carbonate)
George Washington:
BEFORE
George Washington:
AFTER
Chapters 19 and 9.4

62. Naming Acids and Bases Naming Acids: Three Rules Naming Bases
Name of anion ends in –ide
Acid name begins with hydro-
Stem of anion has suffix –ic
Name of anion ends in –ite
Stem of anion has suffix –ous
Name of anion ends in –ate
All three end with the word “acid”
Naming Bases
Named just like ionic compounds – cation + anion
Chapters 19 and 9.4

63. Chapters 19 and 9.4