1. Acids and Bases

2. Electrolytes
Substance that conducts electricity when dissolved in water
Ionizes in water
Examples: aqueous ionic solutions, acids, bases

3. Properties of Acids Acids have a pH less than 7.0
Acids will burn your skin
Dilute acids have a sour taste
Strong acids are good conductors of electricity
Acids react with bases to form water and a salt (neutralization reaction)
Acids react with certain metals to produce hydrogen gas
Metals above H2 will react with acids to produce H2(g)

4. Properties of Bases Bases have a pH greater than 7.0
Bases have a slippery or soapy feeling
Dilute bases have a bitter taste
Strong bases are good conductors of electricity
Bases react with acids to form water and a salt (neutralization reaction)

5. Common Acids and Bases

6. pH Scale Scale ranges from 0-14 Acids = pH less than 7
Bases = pH greater than 7
Neutral = pH = 7

7. pH of common substances

9. Indicators Common indicators are listed in Table M Examples:
What color will thymol blue be in a solution with a pH of 6.5?
What color will methyl orange turn if a solution has a pH of 7.0?
What color will phenolphthalein turn in an acid?

10. Arrhenius Theory Definition of acids and bases
Acids – substances whose water solution contain hydrogen ions (H+), ionize to produce H+
Bases – substances whose water solutions contain hydroxide ions (OH-), ionize to produce OH-
Properties of acids and bases are due to an excess of H+ or OH- ions

11. Hydrogen/Hydronium
H+ cannot exist unbonded in a system; instead it is bonded with a water molecule to make the hydronium ion (H3O+)
H+ and (H3O+) are both used to indicate the presence of an acid

12. Bronsted-Lowry Theory
Acids – donate/lose H+
Bases – gain H+
Example: NH3 + H2O  NH4+ + OH-
H2O is the
NH3 is the

13. Strength of Acids/Bases
Strength is proportional to the degree to which it ionizes in solution
Greater dissociation (more ions), stronger

14. Ionization Constant for Water Kw
For pure water, at 25oC
Kw = 1.0 x = [H+][OH-]
Kw is a constant
therefore [H+][OH-] = 1.0 x 10-14
Examples:
1. Find the [OH-] concentration if [H+] = 10-8
2. Find the [H+] concentration if [OH-] = 10-3
3. Find the [OH-] concentration if [H+] = 10-7
* Exponents add up to -14

15. Hydrogen Ion Concentration (pH)
used for convenience
pH = -log [H+]
Examples:
[H+] = 1.0 x 10-7, pH =
2. [H+] = 1.0 x , pH =
Acids: pH is lower than 7, [H+] is greater than [OH-]
Bases: pH is greater than 7, [H+] is less than [OH-]

18. Hydroxide Ion Concentration (pOH)
pOH = -log [OH-]
Since [H+][OH-] = 1.0 x 10-14
pH + pOH = 14
Examples:
Find the pOH of a solution with a pH = 8. Is this solution acidic or basic?
Find the pH of a solution with a pOH = 12. Is this solution acidic or basic?

19. Concentration and pH
A change of 1 in pH is a tenfold increase in acid or base strength
A pH of 4 is 10 times more acidic than a pH of 5
A pH of 12 is 100 times more basic than a pH of 10

20. Mono/Di/Triprotic Monoprotic acids Diprotic aicds Triprotic acids
produce a single hydrogen ion
Examples: HCl, HBr
Diprotic aicds
produce two hydrogen ions
Examples: H2SO4, H2S
Triprotic acids
produce three hydrogen ions
Examples: H3PO4

21. Neutralization Reactions
An acid and a base react together to form water and an ionic salt
Examples:
HCl + NaOH  H2O + NaCl
H2SO4 + Ca(OH)2  2H2O + CaSO4
HNO3 + Ca(OH)2 

22. Acid-Base Titrations
Process of adding a measured volume of an acid (or a base) of known molarity to a base (or an acid) of unknown molarity until neutralization occurs
Standard Solution – acid or base of known molarity
End Point – point of neutralization
Unknown molarity is calculated using the titration formula

23. Titration MAVA = MBVB equation
MA = Molarity of H+
MB = Molarity of OH-
Moles of acid = moles of base (moles = molarity x liters)

24. Titration Examples
What volume of 1.0M sulfuric acid can be neutralized by 50.0mL of 3.0M sodium hydroxide?
50.0mL of a 0.250M KOH are needed to neutralize 20.0mL of a HCl solution of unknown concentration. What is the concentration of the HCl?