1. Chapter 17 Acid–Base (Proton Transfer) Reactions

2. Arrhenius Acid–Base Theory
An acid ( HCl) is a substance which produces hydrogen ions in water solution. The properties of an acid is the properties of the hydrogen ions.
A base (NaOH) is a substance which produces hydroxide ions in water solution. The properties of a base are the properties of the hydroxide ions.
The net reaction between a strong acid and a strong base is :
 H OH-  H2O
  The Arrhenius concept of acids and base is limited because it applies only to aqueous solutions.

3. Arrhenius Acid–Base Theory
An example of an Arrhenius acid:
Gaseous hydrogen chloride dissolved in water:
HCl(g) H+(aq) + Cl–(aq)
An example of an Arrhenius base:
Solid sodium hydroxide dissolved in water:
NaOH(s) Na+(aq) + OH–(aq)

4. Arrhenius Acid–Base Theory
Properties of an acid must be due to properties of H+.
Properties of a base must be due to properties of OH–.
Thus the cause of the sour taste of acids is the H+ ion, the cause of the bitter taste of bases is the OH– ion.
Other characteristic properties of acids and bases are
also due to the H+ and OH– ions in water solutions.

5. Arrhenius Acid–Base Theory

6. Brønsted–Lowry Acid–Base Theory
An acid is a proton donor.
A base is a proton acceptor.
An acid-base reaction is a proton-transfer reaction in which the proton is transferred from the acid to a base with formation of another acid and base.

7. Brønsted–Lowry Acid–Base Theory
The base formed when the acid has donated a proton is called the conjugate base of the acid.
Acid A ↔ H+ + Conjugate base of acid A
The sign ↔ is used to show that the reaction is reversible.
The stronger the acid, the weaker the conjugate base, and the weaker the acid, the stronger the conjugate base.

8. Brønsted–Lowry Acid–Base Theory
An acid base reaction is a proton transfer reaction in which a proton is transferred from a stronger acid to a stronger base with formation of a weaker acid and weaker base.
Stronger Acid1+ Stronger Base2↔Weaker Acid2+ Weaker Base1
HNO NH ↔ NH NO3-
HCl CH3COO- ↔ HCH3COO Cl-

9. Brønsted–Lowry Acid–Base Theory
H3O+ is called hydronium ion. The conjugate base of acid HNO3 is NO3-

10. Brønsted–Lowry Acid–Base Theory
Water which can behave as a base in one case and an acid in another is said to be amphoteric.

11. Brønsted–Lowry Acid–Base Theory

12. Conjugate Acid–Base Pairs
B HA HB A–
base acid acid base
proton proton proton proton
remover source source remover
Conjugate Acid–Base Pair
Two species that transform into each other by
gain or loss of a proton, H+.
HB+ and B and HA and A– are conjugate acid–base pairs

13. Conjugate Acid–Base Pairs

14. Conjugate Acid–Base Pairs

15. Lewis Acid–Base Theory
Lewis Theory of Acids and Bases
Acid
Electron-pair acceptor.
Base
Electron-pair donor.

16. Lewis Acid–Base Theory

17. Relative Strengths of Acids & Bases

18. Predicting Acid–Base Reactions

19. Autoionization of water
The Water Equilibrium
Autoionization of water

20. Kw is the water constant or equilibrium constant for water
The Water Equilibrium
H2O(l) H+(aq) + OH–(aq)
Kw = [H+] [OH–] = 1.0 × 10–14
Kw is the water constant or equilibrium constant for water
If [H+] = [OH–] = x
Kw = [H+] [OH–] = 1.0 × 10–14 = x2
x = = 10–7 moles/liter

21. The Water Equilibrium For water or water solutions:
If [H+] = [OH–] = 10–7 M,
the solution is neutral.
If [H+] > [OH–],
the solution is acidic.
If [H+] < [OH–],
the solution is basic.

22. The Water Equilibrium Example:
What is the hydrogen ion concentration in a solution of 10–4 M sodium hydroxide in which the hydroxide ion concentration is 10–4 M? Is the solution acidic or basic?
Solution:
GIVEN: [OH–] = 10–4 M WANTED: [H+]
EQUATION: Kw = [H+] [OH–] = 10–14
Since [H+] = 10–10 M< [OH–] = 10–4 M, the solution is basic

23. pH and pOH By definition pH and pOH are given by pH ≡ -log [H3O+]
pOH ≡ -log [OH-]
[H3O+] ≡ antilog(-pH) ≡ 10-pH
[OH-] ≡ antilog(-pOH) ≡ 10-pOH

24. pH and pOH What is the pH of a solution with [H+] = 10–5 M? Solution:
pH = – log [H+] = – log 10–5 = 5
What is the [OH–] of a solution with pOH = 6?
[OH–] = antilog (–pOH) = antilog (–6) = 10–6 M

25. pH and pOH Kw = [H+] [OH–] = 1.0 × 10–14 [H+] [OH–] = 1.0 × 10–14
– log ([H+] [OH–]) = – log (1.0 × 10–14)
– log ([H+] [OH–]) = 14
– log [H+] + (– log [OH–]) = 14
pH + pOH = 14

26. pH and pOH

27. pH and pOH
Example:
The hydrogen ion concentration of a 10–3 M HCl solution is
10–3 M. What are the pH, pOH, and [OH–] of the solution?
Solution:
pH = – log [H+] = – log 10–3 = 3
pH + pOH = 14
pOH = 14 – pH = 14 – 3 = 11
[OH–] = antilog (–pOH) = antilog (–11) = 10–11 M

28. pH and pOH A solution is neutral if [H+] = 10–7 M
A solution is acidic if [H+] > 10–7 M
A solution is basic if [H+] < 10–7 M
Using pH = – log [H+] and pOH = – log [OH–],
A solution is neutral if pH = 7
A solution is acidic if pH < 7
A solution is basic if pH > 7

29. Significant Figures and Logarithms
In a logarithm, the digits to the left of the decimal are not counted as significant figures.
Counting significant figures in a logarithm begins at the decimal point.

30. pH and pOH
Example:
The hydrogen ion concentration of a solution is 2.7 × 10–6 M. What are the pH, pOH, and hydroxide ion concentration?
Solution:
pH = – log [H+] = – log (2.7 × 10–6) = – log (10–6) – log (2.7) =
= 6 – log (2.7) = 5.57 (2 significant figures)
pH + pOH = 14.00
pOH = – pH = – 5.57 = 8.43
[OH–] = antilog(–pOH) = antilog(–8.43) = 10–8.43 M =3.7 × 10–9 M

31. pH and pOH

32. pH and pOH

33. pH and pOH
Measurement of pH

34. Homework
15, 17, 21, 23, 39, 41, 55, 59, 64.