1. Acids & Bases Acids: acids are sour tasting
Arrhenius acid: Any substance that, when dissolved in water, increases the concentration of hydronium ion (H3O+)
Bronsted-Lowry acid: A proton donor
Lewis acid: An electron acceptor
Bases:
bases are bitter tasting and slippery
Arrhenius base: Any substance that, when dissolved in water, increases the concentration of hydroxide ion (OH-)
Bronsted-Lowery base: A proton acceptor
Lewis acid: An electron donor

2. Lone Hydrogen ions do not exist by themselves in solution
Lone Hydrogen ions do not exist by themselves in solution. H+ is always bound to a water molecule to form a hydronium ion

3. General Equation Brønsted-Lowry Theory of Acids & Bases
Conjugate Acid-Base Pairs
General Equation

4. Brønsted-Lowry Theory of Acids & Bases

5. Brønsted-Lowry Theory of Acids & Bases

6. Brønsted-Lowry Theory of Acids & Bases
Notice that water is both an acid & a base = amphoteric
Reversible reaction

7. ELECTROLYTES
Electrolytes are species which conducts electricity when dissolved in water. Acids, Bases, and Salts are all electrolytes.
Salts and strong Acids or Bases form Strong Electrolytes. Salt and strong acids (and bases) are fully dissociated therefore all of the ions present are available to conduct electricity.
HCl(s) + H2O  H3O Cl-
Weak Acids and Weak Bases for Weak Electrolytes. Weaks electrolytes are partially dissociated therefore not all species in solution are ions, some of the molecular form is present. Weak electrolytes have less ions avalible to conduct electricity.
NH3 + H2O  NH OH-

9. Acids & Bases STRONG vs WEAK H2SO4 NaOH HI KOH HBr Ca(OH)2 HCl Sr(OH)2
_ completely ionized _ partially ionized
_ strong electrolyte _ weak electrolyte
_ ionic/very polar bonds _ some covalent bonds
Strong Acids: Strong Bases:
HClO4 LiOH
H2SO4 NaOH
HI KOH
HBr Ca(OH)2
HCl Sr(OH)2
HNO3 Ba(OH)2

10. Acids & Bases One ionizable proton: HCl → H+ + Cl-
Two ionizable protons:
H2SO4 → H+ + HSO4-
HSO4- → H+ + SO42-
Three ionizable protons:
H3PO4 → H+ + H2PO4–
H2PO4- → H+ + HPO42-
HPO42- → H+ + PO4-3
Combined:
H2SO4 → 2H+ + SO42-
Combined:
H3PO4 → 3H+ + PO43-

11. Acids & Bases a. Al(OH)3 + HCl  b. Ba(OH)2 + HC2H3O2 
For the following identify the acid and the base as strong or weak .
a. Al(OH)3 + HCl 
b. Ba(OH) HC2H3O2 
c. KOH + H2SO4 
d. NH3 + H2O 
Weak base
Strong acid
Strong base
Weak acid
Strong acid
Strong base
Weak acid
Weak base

12. Acids & Bases a. Al(OH)3 + HCl  b. Ba(OH)2 + HC2H3O2 
For the following predict the product. To check your answer left click on the mouse. Draw a mechanism detailing the proton movement.
a. Al(OH)3 + HCl 
b. Ba(OH) HC2H3O2 
c. KOH + H2SO4 
d. NH3 + H2O  3
AlCl3 + 3 H2O 2
Ba(C2H3O2)2 + 2 H2O
K2SO4 + 2 H2O 2
NH4+ + OH-

13. Conjugate Acid-Base Pairs

14. Conjugate Acid-Base Pairs

15. Acids & Bases a. Al(OH)3 + 3 HCl  AlCl3 + 3 H2O
For the following Identify the conjugate acid and the conjugate base. The conjugate refers to the acid or base produced in an acid/base reaction. The acid reactant produces its conjugate base (CB).
a. Al(OH) HCl  AlCl3 + 3 H2O
b. Ba(OH)2 + 2 HC2H3O2  Ba(C2H3O2)2 + 2 H2O
c. 2 KOH + H2SO4  K2SO4 + 2 H2O
d. NH3 + H2O  NH4+ + OH-
CB CA
CB CA
CB CA
CA CB

16. TITRATION nacid = nbase MAVA = MBVB
Titration of a strong acid with a strong base
ENDPOINT = POINT OF NEUTRALIZATION = EQUIVALENCE POINT
At the end point for the titration of a strong acid with a strong base, the moles of acid (H+) equals the moles of base (OH-) to produce the neutral species water (H2O). If the mole ratio in the balanced chemical equation is 1:1 then the following equation can be used.
MOLES OF ACID = MOLES OF BASE
nacid = nbase
Since M=n/V
MAVA = MBVB

17. Ma Va = Mb Vb rearranges to Ma = Mb Vb / Va
TITRATION
MAVA = MBVB
1. Suppose mL of hydrochloric acid was required to neutralize mLof 0.52 M NaOH. What is the molarity of the acid?
HCl + NaOH  H2O + NaCl
Ma Va = Mb Vb rearranges to Ma = Mb Vb / Va
so Ma = (0.52 M) (22.50 mL) / (75.00 mL)
= 0.16 M
Now you try:
2. If mL of M HNO3 neutralized mL of KOH, what is the molarity of the base?
Mb = mol/L

18. Molarity and Titration

19. TITRATION nacid = nbase
Titration of a strong acid with a strong base
ENDPOINT = POINT OF NEUTRALIZATION = EQUIVALENCE POINT
At the end point for the titration of a strong acid with a strong base, the moles of acid (H+) equals the moles of base (OH-) to produce the neutral species water (H2O). If the mole ratio in the balanced chemical equation is NOT 1:1 then you must rely on the mole relationship and handle the problem like any other stoichiometry problem.
MOLES OF ACID = MOLES OF BASE
nacid = nbase

20. TITRATION
1. If mL of M LiOH neutralized mL of H2SO4, what is the molarity of the acid?
2 LiOH + H2SO4  Li2SO4 + 2 H2O
First calculate the moles of base:
L LiOH (0.543 mol/1 L) = mol LiOH
Next calculate the moles of acid:
mol LiOH (1 mol H2SO4 / 2 mol LiOH)= mol H2SO4
Last calculate the Molarity:
Ma = n/V = mol H2SO4 / L = M
2. If mL of Ba(OH)2 solution was used to titrate29.26 mL of M HCl, what is the molarity of the barium hydroxide solution?
Mb = mol/L

21. Molarity and Titration
A student finds that mL of a M NaOH solution is required to titrate a mL sample of hydr acid solution. What is the molarity of the acid?
A student finds that mL of a M NaHCO3 solution is required to titrate a mL sample of sulfuric acid solution. What is the molarity of the acid?
The reaction equation is:
H2SO4 + 2 NaHCO3 → Na2SO4 + 2 H2O + 2 CO2

22. Water Equilibrium

23. Equilibrium constant for water
Water Equilibrium
Kw = [H+] [OH-] = 1.0 x 10-14
Equilibrium constant for water
Water or water solutions in which [H+] = [OH-] = 10-7 M are neutral solutions.
A solution in which [H+] > [OH-] is acidic
A solution in which [H+] < [OH-] is basic

24. 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+]

25. The pH scale pH = - log [H3O+] 1 2 3 4 5 6 7 8 9 10 11 12 13 14
The pH scale ranges from 1 to mol/L or from 1 to 14.
pH = - log [H3O+]
acid neutral base

26. 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!

27. PRACTICE PROBLEM #25 10.8 mL 0.0101 g 29.6 mL 5.623 x 10-9 M pH = 4.0
1. How many milliliters of 1.25 M LiOH must be added to neutralize 34.7 mL of M HNO3?
2. What mass of Sr(OH)2 will be required to neutralize mL of M HBr solution?
3. How many mL of M H2SO4 must be added to neutralize 47.9 mL of M KOH?
4. What is the molar concentration of hydronium ion in a solution of pH 8.25?
5. What is the pH of a solution that has a molar concentration of hydronium ion of x 10-5?
6. What is the pOH of a solution that has a molar concentration of hydronium ion of x 10-10?
10.8 mL g
29.6 mL
5.623 x 10-9 M
pH = 4.0
pOH = 4.9

28. GROUP STUDY PROBLEM #25
______1. How many milliliters of 0.75 M KOH must be added to neutralize 50.0 mL of 2.50 M HCl?
______2. What mass of Ca(OH)2 will be required to neutralize 100 mL of M HCl solution?
______3. How many mL of M H2SO4 must be added to neutralize 25.0 mL of M NaOH?
______ 4. What is the molar concentration of hydronium ion in a solution of pH 2.45?
______ 5. What is the pH of a solution that has a molar concentration of hydronium ion of x 10-9?
______ 6. What is the pOH of a solution that has a molar concentration of hydronium ion of x 10-4?