1
INTRODUCTION TO ACIDS
- by definition, acids dissolve in water and increase H+ concentration
Example: HCl (aq)  H+ (aq) + Cl- (aq)
- note ions separate from each other
- acids generally dissolve metals
- concentrated acids dissolve skin and flesh
- stomach acid is a concentrated acid (hydrochloric acid)
- acids taste sour
- vinegar is a weak (acetic acid)
- technically, if a substance is not dissolved in water, it is not an acid. (We won’t make this
distinction yet.)
NOMENCLATURE OF ACIDS
Binary acids HyX
1. Write the prefix hydro2. Write the name of nonmetal anion with –ic suffix
3. Add the word acid
Examples
HBr  hydrobromic acid
HF  hydrofluoric acid
hydroiodic acid  HI
hydrotelluric acid  H2Te
Note: suffix hydro- implies a binary acid with one exception
hydrocyanic acid  HCN
Oxyacids
1. Write the name of the anion
2. Change suffix
a) change – ate to –ic
b) change – ite to –ous
3. Add word acid
Examples:
HClO2  Chlorite anion  Chlorous acid
H2C2O4  Oxalate anion  Oxalic acid
Sulfurous acid  sulfite anion  H2SO3
- note 2 H+ since SO32- has 2- charge
Nitric acid  nitrate anion  HNO3
Phosphoric acid  phosphate anion  H3PO4
2
ACIDS, BASES AND SALTS
Water (H2O) is made of two ions
H+ (aq) – hydrogen ion
OH- (aq) – hydroxide ion
Acids – substances that increase H+ (aq) conc.
Strong acids
- strong electrolytes, i. e., completely dissociate
HCl (aq) = H+ (aq) + Cl- (aq)
- memorize list
HCl, HBr, HI, HNO3, H2SO4, HClO3, HClO4
Weak acids
- weak electrolytes
HF (aq)  H+ (aq) + F- (aq)
HF (aq)  HF (aq) + H+ (aq) + F- (aq)
a lot
a little
a little
H3PO4 (aq)  H+ (aq) + H2PO4- (aq)
Bases – substances that increase OH- concentration
Strong bases
- strong electrolytes
- all alkali metal hydroxides
LiOH, NaOH, KOH, RbOH, CsOH
- some alkaline earth metal hydroxides
Ca(OH)2, Sr(OH)2, Ba(OH)2
Weak bases
- weak electrolytes
- usually increase OH- (aq) conc. “indirectly” by decreasing H+ (aq) conc.
NH3 (aq) + H2O (l)  NH4+ (aq) + OH- (aq)
Neutralization reaction
Acid + Base  Water + Salt
Salt
- ionic compound
- cations and anions remaining after water is made
HNO3 (aq) + KOH (aq)  H2O (l) + KNO3 (aq)
acid
base
water
salt
3
THE AUTOIONIZATION OF WATER
- Water is not a nonelectrolyte, actually it is a very weak electrolyte.
- One out of every 10 million water molecules dissociates into an H+ ion and a OH- ion.
H2O (l)  H+ (aq) + OH- (aq)
Ion product of water
No matter what conditions exist in an aqueous solution, the following expression is
true. (at 25 C)
H  OH  1.0  1014 M 2
 H   - concentration (molarity) of H+
 

 OH  - concentration (molarity) of OH
Ion product is also known as Kw.
K w  H  OH   1.0  10 14 M 2
 
 


Since H  OH  1.0  1014 M 2 is always true
- if we know [H+], we can calculate [OH-]
- if we know [OH-], we can calculate [H+]
Example: If a solution has [H+] = 1.8 x 10-5 M, what is [OH-]?
Example: If a solution has [OH-] = 9.4 x 10-12 M, what is [H+]?
Example: What is the [H+] and [OH-] for pure water?
H2O(l)  H+ (aq) + OH- (aq)
- For every H+ ion formed, an OH- ion is formed.
- Therefore [H+] must equal [OH-]
H OH   1.0  10



14
H   H   1.0  10
 2


M2
7
M
H H   H 


 2

 1.0  10 14 M 2
OH   1.0  10

7
M
Note: When [H+] = [OH-], the amount of acid and base are in equal amounts. The solution is
called neutral.
4
THE HYDRONIUM ION
H+ is simply a proton.
- A proton will hydrogen bond with water
H
H
H
H
O
O
H+
+
H
H3O+ is a hydronium ion.
Current theories of water treat proton as surrounded by 4 or 6 water molecules.
Therefore, H+ may be truly H9O4+ or H13O6+.
H
H
O
H
H
H+
O
O
H
H
O
H
H
H3O+ is a concept, not an actual structure.
Reconsider autoionization of water as
H2O (l) + H2O (l)  H3O+ (aq) + OH- (aq)
pH SCALE
Definition: pH = - log [H+] = -log [H3O+]
pH scale is used strictly for convenience
- convenient to avoid scientific notation.
[H+] = [H3O+] for pure water equals 10-7; therefore, pH = - log(10-7) = - (-7) = 7
All neutral solutions have pH = 7.
If solution is acidic, pH < 7.
If solution is basic, pH > 7.
5
pH scale of common substances
1
2
3
slightly acidic 4
5
6
7
neutral
very acidic
gastric acid
8
lemon juice
9
coke
10
slightly basic 11
12
coffee
very basic 13
milk of magnesia
household ammonia
household bleach
blood
Calculating pH of strong acid and strong base solutions
Strong acids and bases completely dissociate; therefore, [H+] (or [OH-]) equals the
concentration of the solute.
Example: What is the pH of a 0.0142 M solution of HBr?
Since HBr is a strong acid, [H+] = 0.0142 M
Example: What the pH of 1.5 M NaOH solution?
Since NaOH is a strong base, [OH-] = 1.5 M
Calculating [H+] from pH
By definition, pH = - log [H+]
To undo a logarithm (that is, to solve for [H+]), we need to use the logarithm as a power of 10.
 H    10  pH
Example: Calculate the hydrogen ion concentration for a solution with a pH of 4.74.
Example: Calculate the hydroxide ion concentration for a solution with a pH of 10.81.
6
BRØNSTED – LOWRY ACID/BASE THEORY
Acid/base reactions are simply transfer of proton (hydronium ion) from acid to base.
Brønsted – Lowry Acid
- proton donor
- note HCl, H2SO4, HC2H3O2 all donate protons
Brønsted – Lowry Base
- proton acceptor
- note OH- and NH3 accept protons
In Brønsted – Lowry theory, water is an acid or base depending on the circumstances.
HCl (aq) + H2O (l)  H3O+ (aq) + Cl- (aq)
- water accepts proton; therefore, it is a base.
NH3 (aq) + H2O (l)  NH4+ (aq) + OH- (aq)
- water donates proton; therefore, it is a acid.
Conjugate Acids and Bases
Consider a typical acid/base reaction.
NH3 (aq) + H2O (l)  NH4+ (aq) + OH- (aq)
- H2O has donated proton, it is an acid.
- NH3 accepted proton, it is a base.
Consider the reverse reaction.
NH4+ (aq) + OH- (aq)  NH3 (aq) + H2O (l)
- NH4+ has donated proton, it is an acid.
- OH- accepted proton, it is a base.
- An acid is changed into a base and a base is changed into an acid.
- These pairs of acids and bases are called conjugate acid-base pairs.
NH3 (aq) + H2O (l)  NH4+ (aq) + OH- (aq)
base
acid
conjugate
acid
conjugate
base
Another Example:
H2SO4 (aq) + H2O (l)  HSO4- (aq) + H3O+ (aq)
conjugate conjugate
base
acid
base
acid
7
Another Example:
HC2H3O2 + NaOH  H2O + NaC2H3O2
conjugate conjugate
acid
base
acid
base
Relative strengths of acids and bases
The strength of a conjugate base is related to the strength of its acid.
The strength of a conjugate acid is related to the strength of its base.
Consider perchloric acid:
HClO4 (aq) – acid
ClO4- (aq) – conjugate base
HClO4 is a very strong acid, meaning that it fully dissociates.
HClO4 (aq)  H+ (aq) + ClO4- (aq)
- This means that the ClO4- ion has no desire to accept a proton.
- Therefore the perchlorate ion is a very weak base.
Consider aqueous sodium hydroxide:
OH- – base
H2O – conjugate acid
NaOH (aq)  Na+ (aq) + OH- (aq)
- OH- ion has a “healthy appetite” for H+ ions.
- If any additional H+ is added, OH- will grab it to form water.
- OH- is a very strong base.
- The conjugate acid, H2O is a very weak acid.
As the strength of acid increases, the strength of its conjugate base decreases.
As the strength of base increases, the strength of its conjugate acid decreases.
Example: HC2H3O2 is a weaker acid than HF. Which ion will more readily accept protons in
aqueous solution, C2H3O2- or F-?
HC2H3O2 is weaker acid; therefore, C2H3O2- is stronger base.
Thus C2H3O2- will accept a proton more readily.
Example: Pyridine, C5H5N, is a weaker base than ammonia, NH3. Which ion is the stronger
acid, NH4+, C5H5NH+?
C5H5N is weaker base; therefore, C5H5NH+ is stronger acid.
Therefore, C5H5NH+ will donate its proton easier that NH4+.