1. Acids and Bases L

2. Properties of acids Sour Conduct electricity
Strong electrolytes (some)
Weak electrolytes (some)
React with metals to form hydrogen gas
React with hydroxides to form water and salts

3. Properties of bases Bitter Feel slippery
Can be strong or weak electrolytes
React with acids to form water and salts
Change indicators (litmus turns blue)

4. Naming Compound that forms “–OH” in water. Naming ionic compounds
Acids (H)
Bases
If the  ends in “-ide” the name is “Hydro-,” the root of , the suffix “-ic” then “acid”
Ex: HCl = Hydro-chlor-ic acid
HF = Hydro-flor-ic acid
If  ends in “-ate” the name is the root of , and the suffix “-ic”
Ex: H2SO4 = sulferic acid
If  ends in “-ite” the name is the root of , and the suffix “-ous”
Ex: HNO2 = Nitrous acid
(SEE NEXT)
Compound that forms “–OH” in water.
Naming ionic compounds

6. Lets Practice!!
Name an acid or base when given the formula and vice versa.
A. H2SO E. Magnesium Hydroxide
B. HCl F. Phosphoric Acid
C. NaOH G. Sulfurous Acid
D. HNO H. Hydrobromic Acid

7. Acid Base theories Arrhenius
Acid – creates H+/H3O+ as the only positive ion
HCl(aq)  H+(aq) + Cl-(aq)
Base – creates OH- as the only negative ion
NaOH(aq)  Na+(aq) + OH-(aq)
Neutralization – mixing of an acid with a base that produces salt and water
HCl(aq) + NaOH(aq)  H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq)  NaCl(aq) + H2O(l)
H+(aq) + OH-(aq)  H2O(l)

8. NH3(g) + HCl(g)  NH4+ + Cl- NH4Cl(s)
Acid Base Theories
Arrhenius: Limitations
NH3(g) + HCl(g)  NH4+ + Cl- NH4Cl(s)
Since this reaction doesn’t produce hydronium or hydroxide what is it???
Bronsted- Lowery:
Acid – proton donator
Base – proton acceptor
Coordinate covalent bond is formed
One atom provides both electrons for a covalent bond (the base)

9. The great proton grab Bronsted-Lowery
Acids and bases form conjugate acid-base pairs
Whichever base is stronger will exist in higher concentrations at equilibrium.

10. Water Amphiprotic – can be either an acid or base Self ionization
Water falls apart into ions
H2O H+ + OH-
Very small amount
[H+ ] = [OH-] = 1 x 10-7M
A neutral solution.
In water Kw = [H+ ] [OH-] = 1 x 10-14
Kw is called the water dissociation constant.
In water, If you know either [H] or [OH] you can calculate the other

11. Ion Product constant H2O H+ + OH-
Kw is constant in every aqueous solution
[H+] x [OH-] = 1 x 10-14M2
If [H+] > 10-7 then [OH-] < 10-7
If [H+] < 10-7 then [OH-] > 10-7
If we know one, we can determine the other.
If [H+] > 10-7 acidic [OH-] < 10-7
If [H+] < 10-7 basic [OH-] > 10-7

12. Examples: 2. [H+] = 1.0 X 10-10 3. [OH-] = 1.0 X 10-2
Is The solution Acidic, Basic or Neutral? What is the [OH-]/ [H+]?
1. [H+] = 1.0 X 10-5
2. [H+] = 1.0 X 10-10
3. [OH-] = 1.0 X 10-2
4. [H+] = 1.0 X 10-7
5. [OH-] = 1.0 X 10-7

13. Logarithms Powers of ten. A shorthand for big or small numbers.
pH = -log[H+]
in neutral pH = - log(1 x 10-7) = 7
in acidic solution [H+] > 10-7, 100→
pH < -log(10-7)
pH < 7
in base pH > 7
pH > -log(10-7)
[H+] < 10-7, →10-14

14. pH and pOH
pOH = - log [OH-]
[H+] x [OH-] = 1 x 10-14M2
pH + pOH = 14

15. 1 3 5 7 9 11 13 14
Basic
100
10-1
10-3
10-5
10-7
10-9
10-11
10-13
10-14
Acidic
Neutral
[OH-] pH
[H+]
pOH

16. Examples: Calculating pH from [H+] [H+] = 1.0 X 10-6 mol/L
[H+] = .0001M
[OH-] = 1.0 X 10-2 mol/L
[OH-] = 1.0 X mol/L
Using pH to find [H+]
A. 4.0
B. 11.0
C. 8.0

18. Strength vs Concentration
Strong: dissociates fully in water
Example: HCl, HBr, HClO4, HNO3, H2SO4
For every mole of acid placed in solution you get a mole of H+
Monoprotic – can donates only 1 proton
polyprotic – can donate more then 1 proton
Weak: dissociated partially in water
Established equilibrium (Ka or Kb)
Donates one proton but can usually give more

19. Using Ka/b Ka – acid ionization constant
If you know the initial concentration of an acid and the formula you can calculate [H+] given Ka
Same equation as Keq
Ex: CH3COOH + H2O  CH3OOH- + H3O+
What is [H3O+], if 2.0M of acetic acid was dissolved in water and has a Ka of 1.8 x 10-5?

20. Ka = [CH3OOH-][H3O+] [CH3COOH] Ka = 1.8 x 10-5 = [CH3OOH-][H3O+]
CH3COOH + H2O  CH3OOH- + H3O+ What is [H3O+], if 2.0M of acetic acid was dissolved in water and has a Ka of 1.8 x 10-5?
Ka = [CH3OOH-][H3O+]
[CH3COOH]
Ka = 1.8 x 10-5 = [CH3OOH-][H3O+]
1.8 x 10-5 = x(x) 2M
1.8 x 10-5 = X2

21. CH3COOH + H2O  CH3OOH- + H3O+ What is [H3O+], if 2
CH3COOH + H2O  CH3OOH- + H3O+ What is [H3O+], if 2.0M of acetic acid was dissolved in water and has a Ka of 1.8 x 10-5?
1.8 x 10-5 = X2 2M 3.6 x 10-5 = X2 √(3.6 x 10-5) = √(x2) = [H3O+] 6.0 x 10-3 = [H3O+] NOTE: if an acid has a Ka it is WEAK! Strong Acids: HCl, HBr, HI, HNO3, HClO3, HClO4, H2SO4

22. Indicators – dyes that change color in the presence of an acid or base
Hydrangea
Red Cabbage
If grown in acidic soil turns pink
If grown in basic soil turns blue
Chopped and boiled, remaining liquid
Turns pink in acids
Turns green in bases

23. Indicators (common) Active ingredient in laxatives
Litmus Paper
Phenolphthalein
Red Litmus paper
Tests for bases
Acid – no change
Base – turns blue
Blue Litmus paper
Tests for acids
Acid – turns red
Base – no change
Neutral Litmus paper
Tests for both
Active ingredient in laxatives
Acid – Clear and colorless
Bases – Pink
Used in titrations

24. Titrations
Procedure for calculating the pH of an acid/base by reacting it with a known concentration of base/acid
(Macid)(Vacid)(# of H+) = (Mbase)(Vbase)(# of OH-)

25. Buffers or buffer solutions
Resist change in pH caused by the addition of acids or bases
Contains something that reacts with both acids and bases
Mixtures of weak acids and bases
Conjugate pairs
Most common in organic systems
H2CO3/HCO3-
Non-conjugate acid-base pairs
Most common in pharmaceutical (Alka-seltzer)
NH4+/CH3COO-
Amphoteric species
Water
HCO3-, HPO4-2 in blood pH

26. Oxidation and Reduction reactions LEO goes GER
the loss of electrons from atoms of a substance
LEO- Loss of Electrons is Oxidation
Na → Na+ + e-
Na loses an electron and thus, is oxidized.
Oxidizing agent- substance that facilitates oxidation by accepting lost electrons; it is the substance that is reduced.
the gain of electrons by atoms of a substance
GER- Gain of Electrons is Reduction
Cl2 + 2e- → 2Cl-
Cl2 gains electrons and thus, is reduced.
Reducing agent- the substance that facilitates reduction by losing electrons; it is the substance that is oxidized.

27. 2KBr + Cl2 → 2KCl + Br2 Net ionic equation: 2Br - + Cl2 → Br2 + 2Cl –
Half reaction: 2Br - → Br2 + 2e –
2Br– is oxidized to become Br2.
2Br– is the reducing agent
Half reaction: Cl2 + 2e - → 2Cl –
Cl2 is reduced to become 2Cl –
Cl2 is the oxidizing agent