1. Chapter 16
Acids & Bases

2. Acids vs. Bases Section 16.1 Arrhenius model Brønsted-Lowry model Ex:
Oldest model; only applies to compounds that contain H+ or OH- ions
Brønsted-Lowry model
Refers to a compound’s ability to donate or accept an H+ ion
Ex:
HCl NH3 H2O

3. Brønsted-Lowry Acids and Bases Section 16.2
Water as Brønsted-Lowry acid/base:
No such thing as H+ ion in solution (too unstable)
Only H3O+
Proton transfer reactions:

4. Brønsted-Lowry Example
NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)
Which acts as Brønsted-Lowry base? acid?
H2O behaves as an amphoteric compound
Capable of accepting OR donating H+ ions

5. Conjugate Acids/Bases
The reactants and products associated with a proton transfer reaction are known as a conjugate acid-base pair:

6. Relative Strength of Acids and Bases
The strength of an acid depends primarily on the willingness to donate or accept electrons
i.e. strength of conjugate acid/base
Equilibrium for strong acids lies heavily on the side of the deprotonated form and vice versa

7. The Autoionization of Water Section 16.3
Water has a very interesting property due to its amphoterism
Capable of autoionization:

8. Calculating the [H3O+]
An acid is added to water so that the hydrogen ion concentration is 0.25 M. Calculate the hydroxide ion concentration.
See Sample Exercise 16.4 (Pg. 674)

9. The pH Scale Section 16.4
Concentrations of either H3O+ or OH- are typically very small and therefore cumbersome
The pH scale is a logarithmic scale and is much more convenient:
pH = -log[H3O+]
pOH = -log[OH-]
pH + pOH = 14

10. pH of Common Substances

11. Strong Acids and Bases Section 16.5
For a compound to be classified as a strong acid or base it must completely dissociate into ions when placed into aqueous solution
Very weak conjugate bases
No equilibrium
Ex: HCl(g) + H2O(l)  H3O+(aq) + NO3-(aq)
There are 7 strong acids which you will have to remember:
HCl, HBr, HI, HNO3, HClO3, HClO4, H2SO4

12. HF(g) + H2O(l) H3O+(aq) + F-(aq)
Weak Acids Section 16.6
Weak acids have a relatively strong conjugate base (compared to strong acids) and therefore some competition exists for the proton
HF(g) + H2O(l) H3O+(aq) + F-(aq)
As a result, the equilibrium between conjugate acid/base is treated as if it were any other equilibrium

13. Calculating Ka for a Weak Acid
Picric acid, a weak acid, is dissolved in water to prepare a M solution. The pH of the solution was found to be Calculate Ka for picric acid at this temperature.
See Sample Exercise (Pg. 682)

14. Using Ka to Calculate pH
Calculate the pH in a M solution of the weak acid naphthol, for which Ka is 1.7  10-10
See Sample Exercise (Pg. 685)

15. Using Ka to Calculate pH
Calculate the pH of M C6H5COOH. The Ka of C6H5COOH is 6.3 x 10-5.
See Sample Exercise (Pg. 685)

16. pH of Weak Acids
Calculate the pH of a 0.10 M HF solution

17. Percent Ionization
Weak acids, by definition, do not ionize 100% when placed in aqueous solution
It is therefore possible to calculate the extent of ionization (percent ionization)

18. Calculating Percent Ionization
Calculate the percent of HF molecules ionized in a M HF solution.
See Sample Exercise (Pg. 687)

19. Effect of Concentration on Percent Ionization
As the concentration of a weak acid increases, the equilibrium concentration of H3O+(aq) increases
However, the percent ionization decreases as concentration increases

20. Polyprotic Acids
Acids that have more than one acidic proton are referred to as polyprotic acids
H2C2O4 + H2O HC2O4- + H3O+
HC2O4- + H2O C2O42- + H3O+
= 5.6 x 10-2
= 1.6 x 10-5

21. Calculating pH for Polyprotic Acid Solutions
Because the acidic protons of a polyprotic acid are removed in discreet steps, we must calculate the [H3O+] for each step
However, because Ka1 is typically very large compared to Ka2 or Ka3, we may initially assume that the [H3O+] is a result of the first ionization

22. Calculating pH and Concentration of All Species for a Polyprotic Acid
Calculate the pH of a M H2CO3 solution.
For super duper fun, calculate the [CO32-] in solution.
See Sample Exercise (Pg. 689)

23. Weak Bases Section 16.7
There are fewer weak bases compared to the number of weak acids
Equilbria are the same (only involve OH- as product as opposed to H3O+)

24. Calculating pH of Weak Base Solutions
What is the pH of a 1.44 M (concentration of household ammonia) solution of NH3? Kb = 1.8 x 10-5
See Sample Exercise (Pg. 691)

25. Relationship Between Ka and Kb Section 16.8
Examine the following equilibria:
NH4+(aq) NH3(aq) + H+(aq)
NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)

26. Converting Between Ka and Kb
The product of the dissociation constants for any conjugate acid/base pair is always equal to the dissociation constant for H2O:
Ka x Kb = Kw = 1.0 x 10-14
Likewise: pKa + pKb = pKw = 14.00

27. Calculating pH of a Weak Base (Two Approaches)
Calculate the pH of a solution that is M in ammonium ion. Kb = 1.8 x 10-5

28. Acid-Base Properties of Salt Solutions Section 16.9
Hydrolysis reactions
The conjugate base of weak acids are capable of producing hydroxide ions in solution
Raises pH
Conjugate acid of weak bases are capable of producing hydronium ions in solution
Lowers pH

29. Acid-Base Behavior and Chemical Structure Section 16.10
The chemical structure of a compound is what ultimately determines acid/base behavior
Ex: Why does NaOH act as a base whereas CH3OH acts as a weak acid? Why does KH act as a strong base?
Factors that affect?
Polarization of H—X bond (NaH vs. CH4)
Bond strength (HF vs. HCl)
Stability of conjugate base (HNO3 vs. HNO2)

30. Oxyacids
Special attention is paid to oxyacids because the majority of weak acids fall into this category
The strength of these acids is governed by one overriding principle: charge stabilization
Either through electronegativity or resonance

31. Effect of Additional Oxygens
Because oxygen is a very electronegative element, the addition of more oxygen leads to a stronger acid:

32. Carboxylic Acids
A second major class of weak acids are the carboxylic acids
Characterized by the presence of an COOH group