1. CHAPTER 14
ACIDS AND BASES
Properties of Acids
and Bases
1. Acidic Compounds
1. react with metals that are more reactive than hydrogen.
2. react with carbonates such as marble, seashells and limestone.
3. Change litmus paper to red
4. are electrolytes- they conduct electricity when dissolved in water
5. pH is less than 7
6. they taste sour

2. 1. are electrolytes 2. feel slippery 3. taste bitter
2. Basic Compounds
1. are electrolytes
2. feel slippery
3. taste bitter
4. pH is greater than 7
5. Turn litmus paper blue

3. B. Arrhenius Acids
 and Bases
1. Acids
a. monoprotic
b. diprotic
c. triprotic
hydrogen-containing compounds that ionize to yield hydrogen ions (H+)
Acids that contain one ionizable hydrogen
Ex. HNO3
Acids that contain two ionizable hydrogens
Ex. H2SO4
Acids that contain three ionizable hydrogens
Ex. H3PO4

4. 2. Bases
Ionize to yield hydroxide ions (OH-) in solution
Ex. NaOH
C. Brønsted-Lowry
Acids and Bases
 1. Acids
 2. Bases
Hydrogen ion donors
Hydrogen ion acceptors
The Brønsted-Lowry theory of acids and bases includes all of the acids as described by Arrhenius and also some bases that were not included.

5. Example:
NH3(aq) + H2O (l) → NH4+ (aq) + OH- (aq)
B-L base B-L acid ammonium hydroxide
   ion ion
3. conjugate acids and  bases
 a. Conjugate acid
 b. conjugate base
the product formed when a base gains a hydrogen ion
The product that remains when an acid has donated a hydrogen

6.  c. conjugate
 acid-base pair
Consists of two substances related by the loss or gain of a single hydrogen ion.
NH3(aq) + H2O (l) → NH4+ (aq) + OH- (aq)
In this example the NH3 and NH4+ are a conjugate acid-base pair. The H2O and the OH- are also a conjugate acid-base pair.

7. HCl (aq) + H2O (l) → H3O+ + Cl-
In this example HCl is an acid and Cl- is the conjugate base.
H3O+ is the conjugate acid of the base H2O

9. E. Summary of both
 theories
 1. acids
 2. bases
A substance that produces hydrogen ions when dissolved in water. Acids are known as proton (H+) donors.
a substance that produces hydroxide ions (OH-) when place in water. Some bases are proton acceptors.

10. F. Naming rules
 1. acids
To name acids focus on the name of the anion of the acid. The anion may be monatomic such as HCl or polyatomic such as HNO3
1. When the name of the anion ends in
“-ide”, the acid name begins with
“hydro-“. The stem of the anion has the suffix “-ic” and the name is followed by “acid”
Ex. HCl Hydrogen Chloride
When dissolved in water it becomes
Hydrochloric acid

11. 2. When the anion ends in “-ite”, the acid name is the stem of the anion with the suffix “ous” followed by the word “acid”
Ex. HNO2 (NO2 = nitrite)
Name: nitrous acid
3. When the anion ends in “-ate”, the acid name is the stem of the anion with the suffix of “-ic” followed by the word “acid”
Ex. H2CO3 Hydrogen Carbonate
Name: Carbonic Acid

12. Name the following compounds and then name them as acids
H2SO4  HNO3 H2SO3  HNO2
HCl  H3PO4 HC2H3O2
 2. Bases
Naming of bases is just the same as naming ionic compounds
Ex. Ca OH1- ® Ca(OH)2
Calcium Hydroxide
NaOH  KOH  NH3

13. G. Hydrogen Ions
from Water
When water molecules move around and collide they occasionally react and produce
Hydronium ions H3O+ (acid) and hydroxide ions OH- (base)

14. The reaction between two water molecules that produces ions (OH- and H3O+ or just H+)
The concentration of each ion, H+ and OH- in pure water is
1 x 10-7 mol/L
Their concentrations are the same so the solution is said to be neutral. Pure water is a neutral solution.
 1. self ionization of
 water

16. a. Kw
Ion-product constant for water, the product of the concentrations of the hydrogen ions and hydroxide ions
Kw = [H+] X [OH-]
[ ] brackets mean concentration
Kw= [1 x 10-7 M] X [1 x 10-7 M]
Where M stands for (mol/L)
So:
Kw = 1 x M2
If additional H+ ions are added to the solution the equilibrium shifts. The concentrations of hydroxide ions will decrease. This causes the overall Kw to always be 1 x M.

17. b. acidic and basic
 solutions
 1. acidic solution
 2. basic solution
Not all solutions are neutral. To determine if a solution is acidic, basic or neutral, use the following equation.
Kw = [H+] X [OH-]
Acids always have a H+ concentration greater than 1 x 10-7 and the [H+] is greater than [OH-]
Bases have a H+ concentration less than 1 x 10-7 and the [H+] is less than [OH-]
Also known as an alkaline solution.

19. Example: Colas have an
[H+] = 1 x 10-5 M
Is the solution acidic or basic?
The solution is acidic because the [H+] is greater than 1 x 10-7 M
What is the [OH-]?
Kw = [1 x 10-5 M] X [OH-] = 1 x M
[OH-]= 1 x 10-9 M

21.  H. pH
A measure of the hydronium ion (H3O+) concentration in solution
1. Proposed by Søren Sørensen in 1909 while he was brewing beer.
1. calculation pH
It is the negative logarithm of the hydrogen ion concentration.
pH = -log[H+]
If the [H+] = 1 x pH = -log(1 x 10-7) = 7
What is the pH of a solution with an [H+] = 4.5 x 10-5 M?
pH = -log(4.5 x 10-5) = 4.35

22. 2. pH scale
The scale ranges from 0-14.
Which comes from the concentration of H3O+, which ranges from 100 to 10-14M is a high concentration is a low concentration
pH means power of hydrogen
pH of 7 is neutral
pH of 0 is strongly acidic
pH of 14 is strongly basic
pH is measured using
 a. Indicators such as litmus  paper and phenolphthalein.  Which are just dyes.
 b. Instruments such as a pH  meter, which measure the  conductivity of the solution.  

25. a. Strong acid
b. Weak Acid
c. Strong Base
d. Weak Base
Dissociates completely in solution
HCl + H2O à H3O+ + Cl-
(no HCl remains)
Does not completely dissociates, some of the original acid is left.
Example: Acetic acid , only 0.5% reacts in water
dissociates completely in solution to a metal ion and hydroxide ion Example: NaOH
Does not completely dissociate. React to form hydroxide ion and conjugate acid of the base
Example: ammonia, only 0.5% tends to react in water