1. Chapter 10 Acids and Bases in Our Environment
Lecture Presentation
Chapter 10 Acids and Bases in Our Environment
Bradley Sieve
Northern Kentucky University
Highland Heights, KY

2. 10.1 Acids Donate Protons and Bases Accept Them
Taste sour
Cause of taste in vinegar and citric fruits
Essential to the chemical industry
Sulfuric acid is the number one chemical produced in the United States

3. 10.1 Acids Donate Protons and Bases Accept Them
Taste bitter
Feel slippery
React with skin oils to form slippery soaps
Heavily used in industry, with sodium hydroxide being the most common

4. 10.1 Acids Donate Protons and Bases Accept Them
Brønsted–Lowry
Defines acids and bases based on hydrogen ions (H+)
Acids donate an H+
Bases accept H+

5. 10.1 Acids Donate Protons and Bases Accept Them
HCl reacting with water
HCl donates a hydrogen ion and therefore is an acid
Water accepts the hydrogen ion and therefore is a base

6. 10.1 Acids Donate Protons and Bases Accept Them
Ammonia reacting with water
Water donates a hydrogen ion and therefore is an acid
Ammonia accepts the hydrogen ion and therefore is a base

7. 10.1 Acids Donate Protons and Bases Accept Them
Some chemicals can behave as either an acid or a base
Depends on what it is paired with
Water is the main example of this behavior

8. Concept Check
Identify the behavior as an acid or a base of each participant in the reaction
H2PO4– + H3O+  NH3PO4 + H2O

9. Concept Check
In the forward reaction (left to right), H2PO4– gains a hydrogen ion to become H3PO4. In accepting the hydrogen ion, H2PO4– is behaving as a base. It gets the hydrogen ion from the H3O+, which is behaving as an acid.
In the reverse direction, H3PO4 loses a hydrogen ion to become H2PO4– and is thus behaving as an acid. The recipient of the hydrogen ion is the H2O, which is behaving as a base as it transforms to H3O+.

10. 10.1 Acids Donate Protons and Bases Accept Them
Lewis definition
More general than the Brønsted–Lowry definition
Focuses on lone electron pairs
Acids

11. 10.1 Acids Donate Protons and Bases Accept Them
Bases donate a lone pair of electrons
Arrows show the movement of electrons
Water is donating a pair of electrons to form a bond with the H atom

12. 10.1 Acids Donate Protons and Bases Accept Them
Acids accept a lone pair of electrons
HCl is accepting the lone pair from the water

13. Concept Check
How is it possible for carbon dioxide, CO2, to behave as an acid when it has no hydrogen ions to donate?

14. Concept Check
Carbon dioxide behaves as an acid when it accepts the lone pair on the oxygen atom of a water molecule.

15. 10.1 Acids Donate Protons and Bases Accept Them
Salts
Are ionic products of acid-base reactions
NaCl is a common salt but not the only type
HCl + NaOH  NaCl + H2O
HCl + KOH  KCl + H2O

16. 10.1 Acids Donate Protons and Bases Accept Them
Neutralization
A reaction between an acid and a base
Forms salt and water
[insert table 10.1]

17. Concept Check
Is a neutralization reaction best described as a physical change or a chemical change?

18. Concept Check
New chemicals are formed during a neutralization reaction, meaning the reaction is a chemical change.

19. 10.2 Some Acids and Bases Are Stronger Than Others
Strength of an acid or a base is the measure of how much remains after being added to water
A strong acid or base will have little of the original chemical left
A weak acid or base will have a lot of the original chemical left

20. 10.2 Some Acids and Bases Are Stronger Than Others
Hydrochloric acid (HCl) is a strong acid
All of it is converted to ions

21. 10.2 Some Acids and Bases Are Stronger Than Others
Acetic acid (C2H4O2) is a weak acid
Much of the original chemical remains when added to water

22. 10.2 Some Acids and Bases Are Stronger Than Others
The stronger the acid or base, the better it conducts electricity
Pure Water HCl Acetic Acid

23. 10.2 Some Acids and Bases Are Stronger Than Others
Corrosive ability is based on concentration, not the strength of an acid or a base
Highly concentrated solutions tend to be more corrosive
Less concentrated solutions tend to be less corrosive

24. 10.3 Solutions Can Be Acidic, Basic, or Neutral
Amphoteric
A substance that can behave as an acid or a base
Water is a good example

25. 10.3 Solutions Can Be Acidic, Basic, or Neutral
Pure water contains roughly 1 × 10–7 M H+ and OH– ions
The concentration of H3O+ ([H3O+]) multiplied by the concentration of OH– ([OH–]) always equals a constant Kw
Kw = 1 × 10–14

26. Concept Check
In pure water, the hydroxide-ion concentration is 1.0 × 10–7 M. What is the hydronium-ion concentration?

27. Concept Check
1.0 × 10–7 M, because in pure water, [H3O+] = [OH–].

28. Concept Check
What is the concentration of hydronium ions in a solution if the concentration of hydroxide ions is 1.0 × 10–3 M?

29. Concept Check
1.0 × 10–11 M, because [H3O+][OH–] must equal 1.0 × 10–14 = Kw.

30. 10.3 Solutions Can Be Acidic, Basic, or Neutral
Types of Solutions
Neutral is when the solution has equal number of H3O+ and OH– ions
Acidic is when the solution has more H3O+ ions than OH– ions
Basic is when the solution has fewer H3O+ ions than OH– ions

31. 10.3 Solutions Can Be Acidic, Basic, or Neutral

32. Concept Check
How does adding ammonia, NH3, to water make a basic solution when there are no hydroxide ions in the formula for ammonia?

33. 10.3 Solutions Can Be Acidic, Basic, or Neutral
Ammonia indirectly increases the hydroxide-ion concentration by reacting with water:
NH3 + H2O  NH4+ + OH–
This reaction raises the hydroxide-ion concentration, which has the effect of lowering the hydronium-ion concentration. With the hydroxide-ion concentration now higher than the hydronium-ion concentration, the solution is basic.

34. 10.3 Solutions Can Be Acidic, Basic, or Neutral
pH Scale
Used to describe acidity of a solution
pH = –log [H3O+]
Acidic pH < 7
Neutral pH = 7
Basic pH > 7

35. 10.3 Solutions Can Be Acidic, Basic, or Neutral

36. 10.3 Solutions Can Be Acidic, Basic, or Neutral
Two common ways to determine pH

37. 10.4 Buffer Solutions Resist Changes in pH
Resist large changes in pH
Contain two components, each neutralizing an acid or a base
Often prepared by mixing a weak acid and a salt of that weak acid

38. 10.4 Buffer Solutions Resist Changes in pH
A strong acid added to the buffer is neutralized by the base component

39. 10.4 Buffer Solutions Resist Changes in pH
A strong base added to the buffer is neutralized by the acid component

40. Concept Check
Why do most buffer solutions contain two dissolved components?

41. Concept Check
The weak base component neutralizes any incoming acid, and the weak acid component neutralizes any incoming base.

42. 10.4 Buffer Solutions Resist Changes in pH
Blood contains several buffer systems
Maintains pH between 7.35 and 7.45
Outside this range, proteins become denatured
One buffer system in the blood is based on the carbonic acid/sodium bicarbonate system

43. 10.4 Buffer Solutions Resist Changes in pH

44. 10.5 Rainwater Is Acidic
Rainwater is naturally acidic due to CO2
CO2 + H2O  H2CO3
Normal pH falls around 5.6 but ranges from 5 to 7
Acidity can lead to erosion and the formation of caves

45. Acid rain is rain with a pH lower than 5
10.5 Rainwater Is Acidic
Acid rain is rain with a pH lower than 5
Occurs when moisture absorbs pollutants
Formation of sulfuric acid in rainwater
2SO2(g) + O2(g)  2SO3(g)
SO3(g) + H2O(l)  H2SO4(aq)
Corrodes metal, paint, and other exposed substances

46. 10.5 Rainwater Is Acidic
(a), (b) These two photographs show the same obelisk in New York City’s Central Park before and after the effects of acid rain. (c) Many forests downwind from heavily industrialized areas, such as in the northeastern United States and Europe, have been noticeably hard-hit by acid rain.

47. 10.5 Rainwater Is Acidic

48. Concept Check
When sulfuric acid, H2SO4, is added to water, what makes the resulting aqueous solution corrosive?

49. Concept Check
Because H2SO4 is a strong acid, it readily forms hydronium ions when dissolved in water. Hydronium ions are responsible for the corrosive action.

50. 10.6 Carbon Dioxide Acidifies the Oceans
Atmospheric CO2 dissolves in water to form carbonic acid (H2CO3)
Is quickly neutralized in ocean waters by dissolved alkaline minerals
CO2 is not released back into the environment

51. 10.6 Carbon Dioxide Acidifies the Oceans
This trapping of CO2 in the ocean removes basic carbonate compounds from the water
These basic components are used for survival of aquatic life, including coral and shelled organisms
Lowers the pH of the ocean water
0.1 pH over the past 100 years
Has not occurred naturally for 56 million years

52. 10.6 Carbon Dioxide Acidifies the Oceans