1. Acids and Bases Johannes N. Bronsted Thomas M. Lowry
Both independently developed Bronsted-Lowry theory of acids and bases. 1 1 1 1

2. Acids and Bases: A Brief Review
Classical Acids:
Taste sour
Donate H+ (called “H-plus” or “proton”)
Turn litmus red
Generally formed from H-Z, where Z = nonmetal
Classical Bases:
Taste bitter and feel soapy.
Donate OH- (called “O-H-minus” or “hydroxide”)
Turn litmus blue
Generally formed from MOH, where M = metal
Neutralization:
Acid + Base  Salt + water
H-Z + MOH  MZ + HOH
H+ in water is actually
in the form of H3O+,
“hydronium”

3. Acids & Bases
Acids and bases are special kinds of electrolytes. Like all electrolytes they break up into charged particles.
What sets them apart from each other, and other electrolytes is the way that they break up.

4. Practice Identify each of the following as acids/bases/salts: HC2H3O2
K2SO4
KOH
LiOH
HNO3
Acid HC2H3O2 H+ + C2H3O2-
Salt K2SO4  2K+ + SO42-
Base KOH  K OH-
Base LiOH  Li OH-
Acid HNO3  H NO3-

5. Strong/Weak Acids
Acids can be either strong electrolytes or weak electrolytes.
Strong acids (such as HCl) completely break up into their ions:
HCl (aq)  H+(aq) Cl-(aq)
Weak acids (such as HC2H3O2) only partially break up into their ions:
HC2H3O2  H+ (aq) + C2H3O2-(aq)
Weak acids don’t completely break up because they go to equilibrium!

7. Strong/Weak Bases
Bases can be either strong electrolytes or weak electrolytes.
Strong bases (such as NaOH) completely break up into their ions:
NaOH (aq)  Na+(aq) OH-(aq)
Weak bases (such as NH3) only partially break up into their ions:
NH3 (aq) + H2 O  NH4+ (aq) + OH-(aq)
Weak bases don’t completely break up because they go to equilibrium!

8. Bronsted-Lowry Definitions
Bronsted and Lowry felt that this was too limiting, since there are many non-aqueous systems (no water is present).
They came up with the following definitions for acids and bases.
An acid is a proton (H+ ion) donor
A base is a proton acceptor

9. Brønsted-Lowry Acids and Bases
Proton Transfer Reactions
Brønsted-Lowry acid/base definition:
acid donates H+
base accepts H+.
Brønsted-Lowry base does not need to contain OH-.
Consider HCl(aq) + H2O(l)  H3O+(aq) + Cl-(aq):
HCl donates a proton to H2O. Therefore, HCl is an acid.
H2O accepts a proton from HCl. Therefore, H2O is a base.
Water can behave as either an acid or a base.
Amphoteric substances can behave as acids and bases.

10. Relationship between the two models
Bronsted-Lowry Acids/Bases can exist when no water is present
Arrhenius Acids/Bases only exist in water solutions.
All Arrhenius acids and bases can also be classified as Bronsted-Lowry acids and bases.

11. An example
In the reaction below there are no Arrhenius acids or bases present (because no hydronium ions or hydroxide ions are formed).
However, the HCl is acting as a Bronsted-Lowry acid because it is giving a H+ ion to the NH3 (which is acting as a H+ ion acceptor - a base)

12. HCl as a proton donor Consider the following reaction:
Since the HCl gives up a H+ ion to the water it is acting as a Bronsted-Lowry acid. In the process of donating the proton it also forms a hydronium ion, and that makes it an Arrhenius acid as well.

13. Water as a base But what does that make the water molecule?
Since the water molecule is accepting the H+ ion, it is acting as a Bronsted-Lowry base. Since there is no hydroxide ion (OH-) formed, the water is not acting as an Arrhenius base in this reaction.

14. Ammonia as a base Let’s look at another example:
Here the ammonia molecule is accepting a H+ ion and therefore is acting as an Bronsted-Lowry base. However, in the process of reacting with the water it is also forming a hydroxide ion. That makes the ammonia an Arrhenius base as well.

15. But what about the water?
Since the water is giving up a H+ ion, it is acting as a Bronsted-Lowry acid. Since it does not form hydronium ions, it is NOT acting an Arrhenius acid.

16. Amphiprotic
Sometimes a molecule can donate a proton (act as an acid) and sometimes it can accept a proton (act as a base).
Molecules that have this ability to act as both an acid and a base are called amphoteric or amphiprotic.
Water is the most common example of an amphoteric substance.

17. Brønsted-Lowry Acids and Bases
Conjugate Acid-Base Pairs
Whatever is left of the acid after the proton is donated is called its conjugate base.
Similarly, whatever remains of the base after it accepts a proton is called a conjugate acid.
Consider
After H2O (base) gains a proton it is converted into H3O+ (acid). Therefore, H2O and H3O+ are conjugate acid-base pairs.
After HCl (acid) loses its proton it is converted into Cl- (base). Therefore HCl and Cl- are conjugate acid-base pairs.
Conjugate acid-base pairs differ by only one proton.

18. Brønsted-Lowry Acids and Bases
Conjugate Acid-Base Pairs
In each of following Bronsted-Lowry Acid/Base reactions,
Which is acid? base? conjugate acid? conjugate base?
HCl + H2O  H3O Cl-
acid
base
conj. acid
conj. base
NH3 + H2O  NH OH-
base
acid
conj. acid
conj. base
CH3NH2 + H2SO4  CH3NH HSO4-
conj. base
base
acid
conj. acid

19. Reality check
For each of the following reactions identify any Bronsted-Lowry acids and bases.
HNO H2O  H3O NO3-
HNO NH3  NH NO3-
S H2O  HS OH-
HS OH-  S H-OH
HS HCl  H2S Cl-
Are any of the substances above amphoteric?
acid
base
acid
base
base
acid
acid
base
base
acid

20. The Autoionization of Water
In pure water at 25oC, [H3O+][OH-] = 1 x 10-14
and also, [H3O+] = [OH-] = 1 x 10-7
(From now on, for simplification, let’s use the abbreviation:
[H+] = [H3O+]
which means [H]+ = [OH-] = 1 x 10-7
We define pH = -log [H+]
and pOH = -log [OH-]
In pure water at 25oC, pH = pOH = 7.00
pH + pOH = 14
pKw = 14
Acidic solutions have pH < 7.00
Basic solutions have pH > 7.00

21. Brønsted-Lowry Theory of Acids & Bases

22. Brønsted-Lowry Theory of Acids & Bases

23. Conjugate Acid-Base Pairs

24. The pH Scale
Conc. Drano
Battery acid

25. The pH Scale For [H+] = 3.4 x 10-5M, calculate pH, pOH and [OH-]
(b) For [OH-] = 4.4 x 10-3M, calculate pH, pOH, and [H+]
(c) For pH= 8.9, calculate, pOH, [H+], [OH-]
(d) For pOH= 3.2, calculate pH, [H+], [OH-]
The pH meter is the most accurate way to measure pH values of solutions. 25

26. The pH Scale For [H+] = 3.4 x 10-5M, calculate pH, pOH and [OH-]
pH = 4.47, pOH = 9.53, [OH-] = 2.95 x 10-10
(b) For [OH-] = 4.4 x 10-3M, calculate pH, pOH, and [H+]
(c) For pH= 8.9, calculate, pOH, [H+], [OH-]
(d) For pOH= 3.2, calculate pH, [H+], [OH-]
The pH meter is the most accurate way to measure pH values of solutions.

27. The pH Scale For [H+] = 3.4 x 10-5M, calculate pH, pOH and [OH-]
pH = 4.47, pOH = 9.53, [OH-] = 2.95 x 10-10
(b) For [OH-] = 4.4 x 10-3M, calculate pH, pOH, and [H+]
pOH = 2.36, pH = 11.64, [H+] = 2.27 x 10-12
(c) For pH= 8.9, calculate, pOH, [H+], [OH-]
(d) For pOH= 3.2, calculate pH, [H+], [OH-]
The pH meter is the most accurate way to measure pH values of solutions. 27

28. 2. Strong Acids 3. Strong Bases Calculate the pH of 0.2 M HCl .
pH = -log[H+] = -log[0.2] = 0.70
3. Strong Bases
Calculate the pH of 0.2 M NaOH .
pOH = -log[OH-] = -log[0.2] = 0.70
pH = 14 – 0.70 = 13.30
Calculate the pH of 0.2 M Ba(OH)2.
pOH = -log[OH-] = -log[0.4] = 0.40
pH = 13.60

29. The pH Scale For [H+] = 3.4 x 10-5M, calculate pH, pOH and [OH-]
pH = 4.47, pOH = 9.53, [OH-] = 2.95 x 10-10
(b) For [OH-] = 4.4 x 10-3M, calculate pH, pOH, and [H+]
pOH = 2.36, pH = 11.64, [H+] = 2.27 x 10-12
(c) For pH= 8.9, calculate, pOH, [H+], [OH-]
pOH = 5.1, [H+] = 1.26 x 10-9, [OH-] = 7.94 x 10-6
(d) For pOH= 3.2, calculate pH, [H+], [OH-]
The pH meter is the most accurate way to measure pH values of solutions. 29

30. The pH Scale For [H+] = 3.4 x 10-5M, calculate pH, pOH and [OH-]
pH = 4.47, pOH = 9.53, [OH-] = 2.95 x 10-10
(b) For [OH-] = 4.4 x 10-3M, calculate pH, pOH, and [H+]
pOH = 2.36, pH = 11.64, [H+] = 2.27 x 10-12
(c) For pH= 8.9, calculate, pOH, [H+], [OH-]
pOH = 5.1, [H+] = 1.26 x 10-9, [OH-] = 7.94 x 10-6
(d) For pOH= 3.2, calculate pH, [H+], [OH-]
pH = 10.8, [H+] = 1.58 x 10-11, [OH-] = 6.3 x 10-4
The pH meter is the most accurate way to measure pH values of solutions. 30

31. pH A measure of the hydronium ion
The scale for measuring the hydronium ion concentration [H3O+] in any solution must be able to cover a large range. A logarithmic scale covers factors of 10. The “p” in pH stands for log.
A solution with a pH of 1 has [H3O+] of 0.1 mol/L or 10-1
A solution with a pH of 3 has [H3O+] of mol/L or 10-3
A solution with a pH of 7 has [H3O+] of mol/L or 10-7
pH = - log [H3O+]

32. The pH Scale For [H+] = 3.4 x 10-5M, calculate pH, pOH and [OH-]
pH = 4.47, pOH = 9.53, [OH-] = 2.95 x 10-10
(b) For [OH-] = 4.4 x 10-3M, calculate pH, pOH, and [H+]
pOH = 2.36, pH = 11.64, [H+] = 2.27 x 10-12
(c) For pH= 8.9, calculate, pOH, [H+], [OH-]
pOH = 5.1, [H+] = 1.26 x 10-9, [OH-] = 7.94 x 10-6
(d) For pOH= 3.2, calculate pH, [H+], [OH-]
pH = 10.8, [H+] = 1.58 x 10-11, [OH-] = 6.3 x 10-4
The pH meter is the most accurate way to measure pH values of solutions. 32

33. The pH scale
The pH scale ranges from 1 to mol/L or from 1 to 14.
pH = - log [H3O+]
acid neutral base

34. pH + pOH = 14 ; the entire pH range!
Manipulating pH
Algebraic manipulation of:
pH = - log [H3O+]
allows for:
[H3O+] = 10-pH
If pH is a measure of the hydronium ion concentration then the same equations could be used to describe the hydroxide (base) concentration.
[OH-] = 10-pOH pOH = - log [OH-]
thus:
pH + pOH = 14 ; the entire pH range!