This document discusses Lewis acids and bases, strong acids and bases, pH and pOH scales, autoionization of water, and buffer solutions. It provides definitions and examples of Arrhenius acids and bases, Bronsted-Lowry acids and bases, and Lewis acids and bases. It also discusses properties of strong acids and bases, calculating pH and pOH, and how buffer solutions resist changes in pH through acid-base equilibriums.

1. Lewis acid Lewis Bases
by: Dr. Robert D.Craig, Ph.D.

2. Will finish chapter 3…..

3. Acids bases and buffer solutions

4. Strong Acids and Bases
• Give the names and formulas of some strong
acids and bases.
• Explain the pH scale, and convert pH to
concentration (will do later)
• Evaluate solution pH and pOH of strong acids
or bases.

5. In a glass of water

6. Autoionization of Water
• Autoionization of Water
The equilibrium product Kw = [H+] [OH-]
is a constant at a definite temperature due to the
autoionization of water,
H2O = H+ + OH-.
• At 298 K, Kw = 10-14 and the following relationship
in any aqueous solution is obvious,
• ***pOH + pH = 14 at 298 K.

7. Will need this soon
• ***pOH + pH = 14 at 298 K.
pH = -log[H+]

8. We say
• The pH scale is defined as the negative log of
the concentration of H+: pH = -log[H+]
• The pOH scale is defined as the negative log of
the concentration of OH-, [OH-]:
• pOH = -log[OH-] With this scale, calculating
the pOH can be done in the same manner as
the pH scale.

9. .

10. Adapted from
• http://www.science.uwaterloo.ca/~cchieh/cact/c
123/stacids.html
• http://www.chem1.com/acad/webtext/chembon
d/cb03.html
• http://www.epa.gov/acidrain/education/site_stu
dents/phscale.html

11. Arrhenius Acids and Bases
• Arrhenius Acids and Bases
• The Arrhenius definition of acids
and bases is one of the oldest. An
Arrhenius acid is a substance that
when added to water increases the
concentration of H1+ ions present.

12. Arrhenius Acids and Bases
• The chemical formulas of Arrhenius acids are
written with the acidic hydrogens first. An
Arrhenius base is a substance that when
added to water increases the concentration of
OH1- ions present. HCl is an example of an
Arrhenius acid and NaOH is an example of an
Arrhenius base.

13. Arrhenius Acids and Bases
• HCl is an example of an Arrhenius acid and
NaOH is an example of an Arrhenius base.
•
•

14. Arrhenius Acids and Bases
• The H1+ ion produced by an Arrhenius acid is
always associated with a water molecule to
form the hydronium ion, H3O1+(aq).

15. Arrhenius acids
• Arrhenius acids are frequently referred to as
proton donors, hydrogen ion donors, or
hydronium ion donors,

16. Arrhenius Acids and Bases
• To represent the transfer of the H1+ ion to
water to form the hydronium ion, we must
include H2O in the chemical equation for acid
ionization.

17. Arrhenius Acids and Bases

18. Brønsted–Lowry concept
• It follows that, if a compound is to behave as an acid,
donating a proton, there must be a base to accept
the proton. So the Brønsted–Lowry concept can be
defined by the reaction:
• acid + base <-> conjugate base +
conjugate acid.

19. Brønsted-Lowry Style
• Sample Equations written in the Brønsted-
Lowry Style
• A. Reactions that proceed to a large extent:
• HCl + H2O <==> H3O+ + Cl¯
• HCl - this is an acid, because it has a proton
available to be transfered.
• H2O - this is a base, since it gets the proton
that the acid lost.

20. Brønsted–Lowry concept
• The conjugate base is the ion or molecule remaining
after the acid has lost a proton, and the conjugate
acid is the species created when the base accepts the
proton.

21. Now, here comes an interesting idea:
• H3O+ - this is an acid, because it can give a
proton.
• Cl¯ - this is a base, since it has the capacity to
receive a proton.
• Notice that each pair (HCl and Cl¯ as well as
H2O and H3O+ differ by one proton (symbol =
H+). These pairs are called conjugate pairs.

22. .
• The reaction can proceed in either forward or
backward direction; in each case, the acid
donates a proton to the base.

23. Acid base Pairs
2 examples:
CH3COOH + H2O <->CH3COO− + H3O+
H2O + NH3 <-> OH− + NH4+

24. An example
• Which of the following is usually referred to
as strong acid in water solution?
• HF, HNO2, H2CO3, H2S, HSO4-, Cl-, HNO3, HCN
• Answer HNO3
All others are weak acids

25. Water is amphoteric
• Water is amphoteric and can act as an acid or
as a base. In the reaction between acetic acid,
CH3CO2H, and water, H2O, water acts as a
base.
CH3COOH + H2O <->CH3COO− + H3O+

26. conjugate base of acetic acid
• The acetate ion, CH3CO2-, is the conjugate
base of acetic acid and the hydronium ion,
H3O+, is the conjugate acid of the base, water

27. act as an acid
Water can also act as an acid, for instance when it
reacts with ammonia.
The equation given for this reaction is:
H2O + NH3 <-> OH− + NH4 +
• in which H2O donates a proton to NH3

28. reaction as a base!!!
H2O + NH3 <-> OH− + NH4+

29. Strong acid weak acid
• A strong acid, such as hydrochloric acid,
dissociates completely.
• A weak acid, such as acetic acid, may be
partially dissociated; the acid dissociation
constant, pKa, is a quantitative measure of the
strength of the acid.

30. the acid dissociation constant, pKa
• A weak acid, such as acetic acid, may be
partially dissociated; the acid dissociation
constant, pKa, is a quantitative measure of the
strength of the acid.

31. Brønsted–Lowry framework
• A wide range of compounds can be classified
in the Brønsted–Lowry framework: mineral
acids and derivatives such as sulfonates,
phosphonates, etc., carboxylic acids, amines,
carbon acids, and many more

32. Solvent –Not Water????
• Brønsted–Lowry base as the pair of electrons
can be donated to a proton.
• This means that the Brønsted–Lowry concept
is not limited to aqueous solutions.
• Any donor solvent S can act as a proton
acceptor.
•AH + S: <-> A − + SH +

33. donor solvents
• Typical donor solvents used in acid-base
chemistry, such as dimethyl sulfoxide or liquid
ammonia have an oxygen or nitrogen atom
with a lone pair of electrons that can be used
to form a bond with a proton.

34. Do it here Rob
• 87. Most naturally occurring acids are weak
acids. Lactic acid is one example.
• CH3COOH + H2O <->CH3COO− + H3O+

35. Buffer solutions
• If you place some lactic acid in water, it will
ionize to a small extent, and an equilibrium
will be established.
• Suggest an experiment???

36. Who was G.N lewis???

37. G.N. Lewis (1875-1946)
• G.N. Lewis (1875-1946) created the College of
Chemistry at the University of California,
Berkeley, and made it into one of the world’s
most productive centers of chemistry research.
• His other notable work included acid-base theory,
the thermodynamics of solutions, the first
isolation of heavy water (D2O), and the
phosphorescence and magnetic properties of
molecules.

38. Lewis definition

39. Heavy water –nuclear reactors

40. .
• At the time Lewis began developing his ideas
in 1902, it was widely believed that chemical
bonding involved electrostatic attraction
between ion-like entities.

41. Could not explain
• This seemed satisfactory for compounds such
as NaCl that were known to dissociate into
ions when dissolved in water, but it failed to
explain the bonding in non-electrolytes such
as CH4

42. .
• .

43. The ammonium ion-an acid!!!
• The ammonium ion is mildly acidic, reacting
with Brønsted bases to return to the
uncharged ammonia molecule:
• NH4 + + :B- <→ HB + NH
3
• Thus, treatment of concentrated solutions of
ammonium salts with strong base gives
ammonia.

44. .
• When ammonia is dissolved in water, a tiny
amount of it converts to ammonium ions
• H3O+ + NH3 <-> H2O + NH4+

45. depends on the pH
• The degree to which ammonia forms the
ammonium ion depends on the pH of the
solution. If the pH is low, the equilibrium shifts
to the right: more ammonia molecules are
converted into ammonium ions.

46. Just look –last time

47. Strong Acids and Bases
• Acids and bases that are completely ionized
when dissolved in water are called strong
acids and strong bases There are only a few
strong acids and bases, and everyone should
know their names and properties. These acids
are often used in industry and everyday life

48. concentrations of acids and bases
• The concentrations of acids and bases are
often expressed in terms of pH, and as an
educated person, you should have the skill to
convert concentrations into pH and pOH. The
pH is an indication of the hydrogen ion
concentration, [H+].

49. The term Lewis acid
• The term Lewis acid refers to a definition of
acid published by Gilbert N. Lewis in 1923,
specifically: An acid substance is one which
can employ a lone pair from another molecule
in completing the stable group of one of its
own atoms.[1] Thus, H+ is a Lewis acid, since it
can accept a lone pair, completing its stable
form, which requires two electrons

50. Just look – do not write . . .

51. A Lewis base,
• A Lewis base, then, is any species that
donates a pair electrons to a Lewis acid to
form a Lewis adduct. For example, OH− and
NH3 are Lewis bases, because they can donate
a lone pair of electrons

52. Please just look !!!!!!!

53. ammonium ion
• The ammonium ion is mildly acidic, reacting
with Brønsted bases to return to the
uncharged ammonia molecule:
• H3O+ + NH3 <-> H2O + NH4

54. NH4+ + :B- → HB + NH3
• Thus, treatment of concentrated solutions of
ammonium salts with strong base gives
ammonia. When ammonia is dissolved in
water, a tiny amount of it converts to
ammonium ions: (a buffer)

55. How do buffer solutions work?
• A buffer solution has to contain things which
will remove any hydrogen ions or hydroxide
ions that you might add to it - otherwise the
pH will change. Acidic and alkaline buffer
solutions achieve this in different ways.

56. .
• The degree to which ammonia forms the
ammonium ion depends on the pH of the
solution. If the pH is low, the equilibrium shifts to
the right: more ammonia molecules are
converted into ammonium ions.
H3O+ + NH3 <-> H2O + NH4+
NH4 + + :B- → HB + NH
3

57. Alpha curve

58. Buffer solutions
• Buffer solutions achieve their resistance to pH
change because of the presence of an equilibrium
between the acid HA and its conjugate base A-.
• HA <-> H+ + A-
• When some strong acid is added to an
equilibrium mixture of the weak acid and its
conjugate base, the equilibrium is shifted to the
left, in accordance with Le Chatelier's principle

59. These equations are complex

60. Haber process
• Example of l

61. Le Chatelier's Principle
• In 1884 the French chemist and engineer
Henry-Louis Le Chatelier proposed one of the
central concepts of chemical equilibria. Le
Chatelier's principle can be stated as follows:
A change in one of the variables that describe
a system at equilibrium produces a shift in
the position of the equilibrium that
counteracts the effect of this change.

62. Le Chatelier's principle
• Le Chatelier's principle describes what
happens to a system when something
momentarily takes it away from equilibrium.

63. • (1) changing the concentration of one of the
components of the reaction
• (2) changing the pressure on the system
• (3) changing the temperature at which the
reaction is run