1. Acids and Bases
Chapter 19

2. Review
Electrolyte
A substance that conducts an electrical current when melted or in solution
Ionic compounds
Acids and Bases

3. Acid-Base Theories Different definitions of acids and bases Arrhenius
Bronsted-Lowry

4. Arrhenius
Acid
Compounds that ionize to produce hydrogen ions (H+) in aqueous solutions
Examples: HCl, HBr, H2SO4, CH3COOH
*Note: CH3COOH is an organic acid
Acidic H

5. Arrhenius
Base
Compounds that ionize to produce hydroxide ions (OH-) in aqueous solutions
Examples: KOH, NaOH, LiOH
*Note: CH3OH is not a base, it’s an organic alcohol

6. Bronsted-Lowry (alternate)
Acid
Hydrogen ion donor
Examples: HCl, HBr, H3O+
Base
Hydrogen ion acceptor
Examples: H2O, NH3

7. Acid-Base Theories
Bronsted-
Lowry
Arrhenius

8. Properties of Acids & Bases
Acids - Taste Sour
Bases - Taste Bitter, Feel Slippery
Will change color of acid – base indicator
Can be strong or weak electrolytes in an aqueous solution

9. Ionization
Electrolytes will dissociate into ions when dissolved in water
Strong Electrolytes will completely dissociate
Weak Electrolytes will only partially dissociate

10. Ionization HCl(s)  H+(aq) + Cl-(aq) HNO3(s)  H+(aq) + NO3-(aq)
NaOH(s)  Na+(aq) + OH-(aq)
KOH(s)  K+(aq) + OH-(aq)

11. Polyprotic Acids Acids that have more than one H
Can release more than one H+ into solution
Examples:H2SO4, H3PO4
H2SO4(s)  2H+(aq) + SO42-(aq)
Bases can also release more than one OH- into solution
Mg(OH)2(s)  Mg+2(aq) + 2OH-(aq)

12. Ionization of Water Water can be split into 2 ions Ionization of Water
H+ and OH-
Ionization of Water
H2O  H+ + OH-
H2O + H2O  H3O+ + OH-

13. Ionization of Water
H+(aq) = H3O+

15. Neutral Solutions For neutral solutions For all aqueous solutions
[H+] = [OH-]
For all aqueous solutions
[H+] * [OH-] = 1.0 x 10-14
[ ] means concentration

16. Measuring Acidity (Alkalinity)
Traditionally we measure [H+]
pH = -log [H+]
Neutral solution [H+] = 1.0 x 10-7
pH = 7
pOH = -log [OH-]
pH + pOH = 14

17. Acidity Acidic Solutions pH < 7.0 Basic Solutions pH > 7.0
[H+] > 1.0 x 10-7
Basic Solutions pH > 7.0
[H+] < 1.0 x 10-7

19. Measuring pH Litmus paper pH paper pH Meter
Red in acid
Blue in base
pH paper
pH Meter
Acid – Base Indicators (Table M)

20. Table M

22. Changes in pH
pH increases by 1 for every decrease in [H+] by a magnitude of 10
[H+] pH
1.0*10-7 7
1.0*10-8 8
1.0*10-9 9
1.0*10-10 10

23. pH Changes

24. pH Changes

27. Neutralization Acid + Base  Water + Salt HA + BOH  HOH + BA
Double Replacement Reaction
HA + BOH  HOH + BA

28. Neutralization Examples: HCl + NaOH  H2O + NaCl
HNO3 + LiOH  H2O + LiNO3
H2SO KOH  2 H2O + K2SO4

30. Titration
Process in which a volume of solution known concentration is used to determine the concentration of another solution
Usually shown by a color change of an indicator (end point)

31. Titration Example (6M)(1L) = (X)(2L) X =3M NaOH
2 Liters of an unknown conc. of NaOH is titrated with 1 Liter of 6M HCl, what is the concentration of the base?
MAVA = MBVB
(6M)(1L) = (X)(2L)
X =3M NaOH

32. Polyprotic Acids Acids that have more than one H [H2SO4] = 2M
Examples:H2SO4, H3PO4
[H2SO4] = 2M
[H+] = 4M
[H3PO4] = 2M
[H+] = 6M

33. Titration (Reality)
MAVA(#of H’s) = MBVB(#of OH’s)

34. Another Titration Example
500 milliliters of an unknown conc. of NaOH is titrated with 1 Liter of 1M H2SO4, what is the concentration of the base?
MAVA = MBVB
(1M)(1L)(2) = (X)(0.5L)
X =4M NaOH