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