1. Acids and Bases

2. What are acids?
Examples?

3. What are bases?
Examples?

4. 3 different definitions of acids/bases
Arrhenius
Bronsted-Lowry
Lewis
Least broad
Most broad

5. Arrhenius
Acids = a compound that increases the [H+] in aqueous solutions
Bases = a compound that increases the
[OH-] in aqueous solutions

6. Ex of Arrhenius Acid Ex of Arrhenius Base H2O + HCl  H3O+ + Cl-
H2O + NH3  NH4+ + OH-

7. Limitations to Arrhenius’ Definition
Aqueous = in water
Some acids and bases still act as acids or bases even when they aren’t in water.

8. Bronsted-Lowry Acid = molecule or ion that is a PROTON DONOR
Base = molecule or ion that is a PROTON ACCEPTOR
Proton = H+

9. Example NH3 + HCl  NH4+ + Cl-
*Need to memorize that ammonia is NH3 and that it is a base

10. Conjugates Acid base reactions can go in reverse.
Each of the products can be classified as an acid or base as well.
The species that started as an acid becomes the conjugate base and vice versa

11. NH3 + HCl  NH4+ + Cl-
Which product is the conjugate acid? (can donate a H+)
Which product is the conjugate base? (can accept a H+)

12. Amphoteric (Amphiprotic) Compounds
Can act as either an acid or a base (donating or receiving an H+)
WATER is a common example
H2O + CH3COOH  H3O+ + CH3COO-
H2O + NH3  OH- + NH4+

13. Strengths of Acids/Bases
Strong Acids= easily lose protons (100%) Weak acids= some protons are lost Strong Bases= easily accept protons (100%) Weak bases= some accept protons If an acid is strong, the conjugate base is weak, and vice versa

14. Lewis Acid = an electron pair acceptor Base = an electron pair donor
A + :B → A—B
H+ + :NH3 → NH4+
BF3 + F− → BF4−

15. Polyprotic acids Have multiple protons to lose
In excess base (in this case water)
H3PO4 +H2O  H2PO4 - +H3O+
H2PO4 - + H2O  HPO H3O+
HPO H2O  PO H3O+

16. Polyprotic acids continued…
Each time that a polyprotic loses an H+ it becomes harder to lose. Why?
Therefore which acid in a polyprotic is the most acidic?

17. Prefixes: Di, Mono, Poly Monoprotic = only has one H+ to lose
Diprotic = has two H+ to lose (H2SO4)
Polyprotic = has multiple (poly) H+ to lose

18. Homework
Chapter 16: #1,15,18,20,22,24,27,28

19. Autoionization of Water
Pure water self ionizes to a small extent
H2O H+ + OH-
H+  H3O+ (Hydronium ion) (Attaches onto a water molecule)
DYNAMIC equilibrium- no single molecule stays ionized for long

20. The amount it ionizes is very small.
In pure water [H3O+ ] = [OH-] = 1.00x10-7 M
K expression: (Kw stands for water ionization constant)
Kw = [H3O+ ] * [OH-]
K = [1.00x10-7 M ][1.00x10-7 M ] = 1.00x10-14 M (at 25 deg. Celsius)

21. Kw can be used to calculate hydronium ion or hydroxide ion concentrations at any time. Together their product is always 1.00x10-14M If [H3O+] > [OH-] then the solution is acidic If [OH-] > [H3O+] then the solution is basic If they are equal (and therefore both 1.00 x 10-7 M) the solution is neutral

22. Example
Determine the hydronium ion concentration if a solution has a hydroxide concentraion of M. Is this an acidic, basic, or neutral solution?

23. pH
Hydronium power or potential Negative Log scale of [H3O+] pure water has a pH of 7 because -log(1.00x10-7) = 7 Higher pH = lower concentration of Hydronium Lower pH = higher concentration of H3O+

25. pOH is the same log scale, but for OH- pOH + pH = 14 (because [OH-]
pOH is the same log scale, but for OH- pOH + pH = 14 (because [OH-]*[H3O+]=1.00x10-14)

26. Helpful box pH Convert using pH+pOH=14 pOH
Convert using: pH= -log[H3O+] Convert using: pOH= -log[OH-]
Or 10-pH = [H3O+] Or 10-pOH = [OH-]
[H3O+] Convert using Kw [OH-]

27. Examples What is the pH of a 0.040 M HCl solution?
What is the pH of a 0.005M H2SO4 solution?
What is the pH of a 0.008M Ca(OH)2 solution?

28. Homework
31,33,40,46,50