1. Chapter 4: Types of Chemical Reactions
Goals:
To be able to predict chemical reactivity.
To know how to synthesize specific compounds.

2. Types of Reactions Acid-Base: proton-transfer
Oxidation-Reduction: electron-transfer
Precipitation: formation of insoluble salts
Gas Forming
Organic:
Substitution
Addition
Elimination

3. Reactions in Aqueous Solution
Unless mentioned, all reactions studied this and next week occur in aqueous solution.

4. Electrolytes
Strong Electrolytes: solute breaks apart to give ions in solution.
NaCl  Na+ + Cl-
Weak Electrolytes: solute partially breaks apart to give ions.
CH3CO2H  CH3CO2- + H happens less than 5%
Nonelectrolytes: no ions formed.
CH3CH2OH

5. Brønsted-Lowery Acid-Base Definitions
An acid is a substance that donates a proton (H+) to a base
A base is a substance that accepts a proton (H+) from an acid

6. Brønsted-Lowery Definitions
acid: donates a proton (H+) to a base
base: accepts a proton (H+) from an acid
Acid-base reactions can be reversible: reactants  products or products  reactants

7. Brønsted-Lowery Definitions
An acid is a substance that donates a proton (H+) to a base
A base is a substance that accepts a proton (H+) from an acid
Acid-base reactions can be reversible: reactants  products or products  reactants

8. Important Acids and Bases
Strong Acids:
HCl hydrochloric
HBr hydrobromic
HI hydroiodic
HNO3 nitric
H2SO4 sulfuric
HClO4 perchloric
Weak Acid:
CH3CO2H acetic
Any other acids are WEAK
Strong Bases:
LiOH lithium hydroxide
NaOH sodium hydroxide
KOH potassium hydroxide
Ca(OH)2 calcium hydroxide
Ba(OH)2 barium hydroxide
Weak Base:
NH3 ammonia

9. STRONG acids in water: 100% of acid molecules form ions:
HCl(aq) + H2O(l)  H3O+(aq) + Cl-(aq)
H3O+ is
hydronium ion

10. WEAK acids in water, ~5% or less of acid molecules form ions (acetic, H3PO4, H2CO3)

11. Polyprotic Acids multiple acidic H atoms
H2SO4  H+ + HSO4-
HSO4-  H+ + SO42-
Not all H’s are acidic:
CH3CO2H

12. If H3PO4 reacts as an acid, which of the following can it not make?

13. If C2O42- reacts in an acid-base reaction, which of the following can it not make?
1. H2C2O4
2. HC2O4-
3. 2 CO2

14. Acid-Base Reactions Strong Acid + Strong Base
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l) acid base “salt” water

15. Acid-Base Reactions Diprotic Acids or Bases H2SO4(aq) + NaOH(aq) 
H2SO4(aq) Ba(OH)2(aq) 
HCl(aq) + Ba(OH)2(aq) 

16. Acid-Base Reactions Strong Acid + Weak Base
HCl(aq) + NH3(aq)  NH4Cl(aq)

17. Acid-Base Reactions Weak Acid + Strong Base
HCN(aq) + NaOH(aq)  NaCN(aq) + H2O(l) acid base “salt” water

18. Net Ionic Equations HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
What really happens:
H+(aq) + OH-(aq)  H2O(l)
Sodium ion and chloride ion are “spectator ions”

19. Reactions involving weak bases
HCl(aq) + NH3(aq)  NH4+(aq) + Cl-(aq)
Net-Ionic Equation:
NH3(aq) + H+(aq)  NH4+(aq)

20. CH3CO2H(aq) + NaOH(aq)  1. CH3CO2H2+(aq) + NaO(aq)
2. CH3CO2-(aq) + H2O(l) + Na+(aq)
3. CH4(g) + CO2(g) + H2O(l)

21. HCN(aq) + NH3(aq)  1. NH4+(aq) + CN-(aq) 2. H2CN+(aq) + NH2-(aq)
3. C2N2(s) + 3 H2(g)

22. Solution Concentration: Molarity
Molarity = moles solute per liter of solution
0.30 mol NH3 dissolved in L Concentration =
Written like: [NH3] = 0.60 M

23. pH Scale
In pure water, a few molecules ionize to form H3O+ and OH– H2O + H2O  OH– + H3O+
In acidic and basic solutions, these concentrations are not equal acidic: [H3O+] > [OH–] basic: [OH–] > [H3O+] neutral: [H3O+] = [OH–]

24. pH Scale Measure how much H3O+ is in a solution using pH
pH < 7.0 = acidic
pH > 7.0 = basic
pH = 7.0 = neutral
Measure of H3O+ and OH– concentration (moles per liter) in a solution
As acidity increases, pH decreases

25. pH Scale
The pH scale is logarithmic: log(102) = log(101) = log(100) = –1 log(10–1) = – –2 log(10–2) = –2
pH = –log [H3O+]
pH if [H3O+] = 10–5? 10–9? Acidic or basic?
pH if [H3O+] = M?

26. Finding [H3O+] from pH
[H3O+] = 10-pH
What is [H3O+] if pH = 8.9?

27. pH: Quantitative Measure of Acidity
Acidity is related to concentration of H+ (or H3O+)
pH = -log[H3O+]