Acid and base

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 NH4
+ + 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 10-14.

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 10-14.

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 < 7.0 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) CaCl2 + 2H2O(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.