1. Acids, Bases, and Salts
Chapter 19

2. Acid-Base Theories Essential Question:
What are the properties of acids and bases, and what distinguishes the Arrhenius, Bronsted-Lowry and Lewis theories of acids and bases?

3. Properties of Acids Acids taste tart or sour.
Aqueous solutions of acids are electrolytes.
Acids change the color of acid-base indicators.
Acids react with metals to produce hydrogen gas.
Acids react with bases to form water and a salt.

4. Properties of Bases Bases have a bitter taste.
Bases have a slippery feel.
Aqueous solutions of bases are electrolytes.
Bases change the color of acid-base indicators.
Bases react with acids to form water and a salt.

5. Three Major Theories
Arrhenius
Bronsted-Lowry
Lewis

6. Arrhenius Acids and Bases
Acids are hydrogen-containing compounds that ionize to produce H+ ions in aqueous solution.
Bases are hydroxide-containing compounds that ionize to produce OH– ions in aqueous solution.

8. Mono- Di- and Tri-protic Acids
HNO3 is a monoprotic acid.
H2SO4 is a diprotic acid.
H3PO4 is a triprotic acid.
Not all substances that contain hydrogen are acids.
Not all hydrogens in acids are necessarily released as H+ ions.

9. Hydrochloric Acid

10. Ethanoic Acid (Acetic Acid)

11. Arrhenius Bases The most commonly known is NaOH (lye).
Another common base is KOH

13. Bronsted-Lowry Acids and Bases
Defines an acid as a hydrogen-ion donor.
Defines a base as a hydrogen-ion acceptor.

14. Why Ammonia is a Base

15. Conjugate Acids and Bases
When a substance donates a hydrogen ion, what remains has the ability to accept it back.
When a substance accepts a hydrogen ion, what remains has the ability to donate the ion.

16. Conjugate Acids and Bases
A conjugate acid is formed when a base gains a hydrogen ion.
A conjugate base is remains when an acid has donated a hydrogen ion.
A conjugate acid-base pair consists of two substances related by the loss or gain of H+.

17. Consider Ammonia in Water
NH3 + H2O NH4+ + OH–
Base Acid Conjugate Conjugate
Acid Base

19. Hydronium ion ( H3O+ )
A water molecule that gains a hydrogen ion becomes a positively charged hydronium ion.

20. Amphoteric Sometimes water accepts a hydrogen. HCl + H2O H3O+ + Cl–
Other times, water donates a hydrogen.
NH3 + H2O NH OH–
A substance that can act as either an acid or a base is said to be amphoteric.

21. Lewis Acids and Bases
A Lewis acid accepts a pair of electrons during a reaction.
A Lewis base donates a pair of electrons during a reaction.

22. Hydrogen Ions and Acidity
Essential Question:
How are [H+] and [OH–] related and how do they relate to acidity?

23. Hydrogen Ions from Water
The reaction in which water molecules produce ions is called the self-ionization of water.
H2O (l) H+ (aq) + OH– (aq)
In water, H+ ions are always joined to water molecules to form H3O+ hydronium ions.

24. Self-Ionization of Water

25. Self-Ionization of Water
This happens to a very small extent.
In pure water, the concentration of H+ and OH– are equal.
[H+] and [OH–] both equal 1.0 x 10-7 M.
This is called a neutral solution.

26. Ion Product Constant for Water
In any aqueous solution, when [H+] increases, [OH–] decreases.
When [H+] decreases, [OH–] increases.
The product of the hydrogen-ion concentration and the hydroxide ion concentration always equals 1.0 x

27. Ion Product Constant for Water
Kw = [H+] x [OH–] = 1.0 x 10-14
Remember, as [H+] goes up, [OH–] goes down.
But the product of [H+] x [OH–] will always be 1.0 x

28. Acidic vs Basic
An acidic solution is one in which the [H+] is greater than the [OH–].
A basic solution is one in which the [H+] is less than the [OH–].
Basic solutions are also called alkaline solutions.

32. The pH Concept
Expressing hydrogen concentration in molarity can be cumbersome.
A more popular method is the pH scale.
The pH scale ranges from 0 to 14, with the most acidic = 0 and the most basic = 14.

33. Calculating pH
pH = the negative of the logarithm of the hydrogen ion concentration.
pH = -log [H+]

34. Calculating pOH pOH is the red-headed stepchild…
pOH is the negative logarithm of the [OH–].
pOH = -log [OH–]

35. pH and pOH
You can calculate the pH or the pOH of a solution using the log function key on a calculator.

36. pH and Acidity
Acidic solution: pH < and [H+] is greater than 1 × 10−7M
Neutral solution: pH = 7.0 and [H+] is equal to 1 × 10 −7M
Basic solution: pH > 7.0 and [H+] is less than 1 × 10 −7M

39. pH and Significant Figures
Express [H+] and [OH–] in scientific notation.
Express pH and pOH with the same number of digits to the right of the decimal place as the number of significant digits in the scientific notation.

41. Sample Problem 19.2

42. The Solution

43. Measuring pH

44. Measuring pH Two common methods are used:
Acid base indicators which change color
pH meters which measure electrical conductivity

45. Acid-Base Indicators
Acid-Base indicators have different colors based upon acidity.
For each indicator, the change takes place over a relatively narrow range of about 2 pH units.

48. Effects of Acidity on Plant Color

49. Strengths of Acids and Bases
Essential Question:
How does the value of an acid dissociation constant relate to the strength of the acid, and how are those values calculated?

50. Strong and Weak Acids
Strong acids are completely ionized in aqueous solution.
HCl + H2O H3O+ + Cl– (100% ionized)
Weak Acids only slightly ionized in aqueous solution.
CH3COOH + H2O H3O + CH3COO–
(partially ionized)

52. Acid Dissociation – Strong

53. Acid Dissociation – Weak

54. Acid Dissociation – Very Weak

55. Acid Dissociation Constants
Takes the same form as the equilibrium-constant expression from a balance chemical equation.
For ethanoic acid, for instance, the acid dissociation constant is calculated as follows:
[H3O+] x [CH3COO–]
Ka =
[CH3COOH] x [H2O]
These are sometimes called ionization constants.

56. Acid Dissociation Constants
Weak acids have small Ka values.
Strong acids have large Ka values.
The Ka value for HCl (aq) is ∞ (infinite).
Why do you think this is the case?

58. Base Dissociation Constants
Strong bases and weak bases refer to the degree of dissociation just like acids.
For sodium hydroxide, for instance, the base dissociation constant is calculated as follows:
[Na+] x [OH–]
Kb =
[NaOH] x [H2O]

59. Neutralization Reactions
Essential Question:
What are the products of the reaction of an acid and a base when the endpoint of a titration is reach?

60. Acid—Base Reactions
A mixture of a strong acid with an equal amount of a strong base results in a neutral solution.
These reactions are called neutralization reactions.

61. Neutralization Reactions
Consider these examples:
HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)
H2SO4(aq) + 2KOH(aq) K2SO4(aq) + H2O(l)
What would the net ionic equations for each of these reactions be?

62. Titration Acids and bases sometimes react 1:1
HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)
However, the ratio can vary
H2SO4(aq) + NaOH(aq) Na2SO4(aq) + 2H2O(aq)
2HCl(aq) + Ca(OH)2(aq) CaCl H2O(l)

63. Titration
When an acid and a base are mixed, the equivalence point is when the number of moles of hydrogen ions equals the number of moles of hydroxide ions.

64. Titration
You can determine the concentration of an acid in a solution by performing a neutralization reaction.
You must select an appropriate acid-base indicator.
Phenolphthalein turns from colorless to pink as the pH changes from acidic to basic.

65. Titrations
A measured volume of an acid solution of unknown concentration is added to a flask.
Several drops of the indicator are added to the solution.
Measured volumes of a base of known concentration are mixed into the acid until the indicator just barely changes color.