1. Unit 12 – Organic, Nuclear, Oxidation-Reduction and Acids and Bases

2. Acids and Bases
Groups of chemical compounds that share certain characteristics
common substances you use everyday
Examples:
soap juice
drain cleaners soda
vinegar antacids

3. Acid Properties 1. sour taste when dissolved in water
do not taste to test chemicals
many are corrosive and poisonous
2. turns pH paper red
3. react with metals to make H2(g)
metals above H2 in activity series
single replacement reaction
Ba(s) + H2SO4(aq)  BaSO4(aq) + H2(g)

4. Acid Properties 4. react with bases to make salt and water
acid is neutralized when added to an equal number of moles of base
produces an ionic compound (also called “salt”) and water
HBr(aq) + KOH(aq)  H2O(l) + KBr(aq)
5. conduct electricity
strong acids conduct very well
weak acids conduct a small amount

5. Common Acids Sulfuric acid: H2SO4 Nitric acid: HNO3
used in car batteries
Nitric acid: HNO3
used in explosives
Phosphoric acid: H3PO4
used as flavoring in soft drinks
Hydrochloric acid: HCl
stomach acid
Acetic acid: HCH3COO or HC2H3O2
vinegar contains it

6. Base Properties 1. taste bitter 2. turn pH paper blue
most are caustic (cause burns) so don’t taste to test
2. turn pH paper blue
3. feel slippery when dissolved in water
4. react with acids to make salt and water
5. conduct electricity (strong vs. weak)

7. Naming of Acids (AGAIN )
Binary acids
Contains 2 different elements: H and another
Always has “hydro-” prefix
Root of other element’s name
Ending “-ic”
Examples: HI, H2S, HBr

8. Naming Acids Ternary Acids - Oxyacids
Contains 3 different elements: H, O, and another
No prefix
Name of polyatomic ion
Ending “–ic” for “-ate” and “–ous” for “-ite”
Examples: HClO4, H3PO4, HNO2

9. Types of Acids Monoprotic acids Polyprotic acids- diprotic, triprotic
Only have one acidic proton
HNO2, HBr, HCH3COO
Polyprotic acids- diprotic, triprotic
Contains more than one acidic proton
H3PO4, H2SO4, H2Te
Organic acids
Contain COOH (carboxyl) group
Weak acids

10. Strengths of acids
Depends on polarity and strength of the bond holding the H to the rest of the molecule
Strong acids
Ionize completely
Strong electrolyte
Would these have a strong or weak bond?
Would these have a polar or nonpolar bond?

11. Strengths of Acids Strong Acids Weak Acids
HCl, HBr, HI, HClO4, H2SO4, HNO3
Memorize that list!
Weak Acids
Only ionize partially
Weak electrolytes
Can assume any acid not on that list is weak

12. Strengths of Bases Depends on the amount of dissociation and structure
Strong Bases
Strong electrolytes
All are Group I and II metal hydroxides
Dissociate completely in water
Examples: KOH, Ba(OH)2
Weak Bases
Weak electrolytes
Examples: NH3, C6H5NH2

13. Definitions of Acids and Bases
Arrhenius
Most specific/exclusive definition
Created by Svante Arrhenius, Swedish
Acid : compound that creates H+ in an aqueous solution
Base : compound that creates OH- in an aqueous solution
HNO3  H+ + NO3-
NaOH  Na+ + OH-

14. Definitions of Acids and Bases
Bronsted-Lowry
A bit more general and more common
Created by two scientists around the same time (1923)
Acid: Molecule or ion that is a H+ donor
Base: Molecule or ion that is a H+ acceptor
HCl + H2O  H3O+ + Cl-
Does not include strong bases

15. Acids or Bases? C5H5N Mg(OH)2 HF NH3 H2CO3 H2SO4 KOH CH3NH2 HC2H3O2 HI
base (weak)
Mg(OH)2
base (strong) HF
acid (weak)
NH3
H2CO3
H2SO4
acid (strong)
KOH
base (strong)
CH3NH2
base (weak)
HC2H3O2
acid (weak) HI

16. Acid-Base Reactions Amphoteric Substances
Can act as either acid or base depending on what they are mixed with
Example: Water, HSO4-

17. pH Scale
Strength of an acid or a base solution refers to the concentration [ ] of H+ ions in solution  numbers can be very small  developed the pH scale
Scale from 0-14 (values beyond the scale do exist for very strong acids or bases)
pH increases as [H+] decreases (inversely or indirectly related)

18. pH Scale
pH < 7: acid
pH > 7: base
pH = 7: neutral

19. Logarithmic Scale
each whole pH value below 7 is ten times more acidic than the next higher value. For example, pH 4 is ten times more acidic than pH 5 and 100 times (10 times 10) more acidic than pH 6.

21. Math & pH scale
Can use the equation below to find the [H+] or the pH of a solution
pH=-log [H+]
H+ = H3O+  will see both in different sources

22. Example #1
What is the pH of a solution if the [H+] is: 3.4 x 10-5 M? Is the solution acidic or basic?
pH = -log [H+]
pH = -log 3.4 x 10-5)
pH = 4.47  acidic solution
(No units on pH values)

23. Example #2
What is the hydrogen ion concentration of an aqueous solution that has a pH of 4.0?
pH = -log [H+]
4.0 = -log [H+]
-4.0 = log [H+]
antilog (-4.0) = [H+] or 10x key
1.0 x 10-4 M = [H+]