Acid-base

Acids and Bases

Ways to Define Acids and Bases

A. Operational Definition (defined by how it works)
	1. Electrolytes – conduct electricity when dissolved in water.
	2. In dilute solutions	- acids taste sour - bases taste bitter
	3. Bases(aq) feel slippery (like soap).
	4. Acids turn litmus red and Phenolphthalein colorless.
	5. Bases turn litmus blue and Phenolphthalein pink.
	6. Neutralization reaction: Acid + Base Salt + Water (don't forget that a salt is any ionic substance, not just the stuff you put on your fries.)

B. Arrhenius definition – (Swedish chemist)
1. Arrhenius Acid: a substance that releases H+ ions as the only positive ion in aqueous solution.
2. Arrhenius Base: a substance that releases OH- ions as the only negative ion in aqueous solution.
Example: NaOH + HCl NaCl + HOH (H2O)
Note: an Arrhenius acid base pair will always undergo a neutralization reaction to make a salt and water.

C. Bronsted – Lowry Definition
	1. Bronsted–Lowry Acid – any species that can donate a proton (H+ ions) in solution PROTON DONOR
	2. Bronsted – Lowry Base – any species that accepts a proton (H+ ions) in solution PROTON ACCEPTOR
Example: NH3 + HCl NH4+ + Cl-
Note: a Bronsted-Lowry acid base pair can undergo a reversible reaction to produce a conjugate acid base pair.

II.
Acid Base Equilibrium

A. Conjugate Pairs:
1. In an acid base equilibrium there is an acid and a base in the forward reaction and an acid and a base in the reverse reaction.
2. The acid on one side reacts to become the base on the other side.
HCl + NH3NH4+ + Cl-
If we watch the movement of the proton (H+) we can identify acids and bases in the reaction.

	Acid 1 (Donor)	Base 1 (Acceptor)	Acid 2 (Donor)	Base 2 (Acceptor)
Acid 1 (HCl) and Base 2 (Cl-) are conjugate pairs
Acid 2 (NH4+) and Base 1 (NH3) are also conjugate pairs
The only difference between and acid and its conjugate base is a proton (H+)

B. Ionization Constants
eq. HCl (aq) H30+ (aq) + Cl- (aq)
Ka =	[H30+] [Cl-]
	[HCl]
The higher the Ka the stronger the acid
eq. NaOH (aq) Na (aq) + OH- (aq)
Kb =	[Na] [OH-]
	[NaOH]
The higher the Kb the stronger the base
eq. H20 (l) + H20 (l) (aq) H30+ (aq) + OH- (aq)
[H30+] [OH-] = Kw = 1.0 x 10 ^–14
Since water is equal parts acid and base
[H30+] = [OH-] therefore
Kw = X·X = X^2
X^2 = 1.0 x 10^–14 ....and.... X = 1.0 x 10^ –7

C. The Relative Strengths of Acids and Bases

The strength of an acid or base is determined by how well it
dissociates (breaks-up) in solution. The more it breaks up the
more ions produced. The more H+ ions produced the stronger
the acid. The more OH- ions produced the stronger the base.

1. Based on [H30+]

[H30+] = 1.0 x 10 ^–x
X = pH

When [H30+] is large X is small (1)

When [H30+] is in the middle X is average (7)

When [H30+] is small X is large (14)

2. pH Scale

	0-2	Strong acid
	3-5	Moderate acid
	6	Weak acid
	7	Neutral
	8	Weak base
	9-11	Moderate base
	12-14	Strong base

3. Based on [OH-]

[OH-] = 1.0 x 10 –x X = pOH

When [OH-] is large X is small (1)

When [OH-] is in the middle X is average (7)

When [OH-] is small X is large (14)

4. Relationship between pH and pOH

pH + pOH = 14

5. Strengths of Conjugate Pairs

There is an inverse relationship.
If the acid is strong its conjugate base is weak.

6. Strength of products of Neutralization reaction
Not all neutralizations result in a neutral solution.

Strong Acid + Strong Base = Neutral Solution
Strong Acid + Weak Base = Acidic Solution
Weak Acid + Weak Base = Neutral Solution
Weak Acid + Strong Base = Basic Solution

III.
Acid-Base Titrations

A. An acid and a base neutralize to form a salt and water.

H+ + OH- H20

B. In order to have complete neutralization the moles (n) of acid must be equal to the moles (n) of base.

n(acid) = n (base)

C. When dealing with solutions we talk about molarity(M) instead of moles(n)

Molarity (M)/Volume = moles (n)

Therefore:

n = M x V for the acid
and
n = M x V for the base

Therefore:

M(acid) x V(acid) = M(base) x V(base)

When solving these problems you will be given three of the four variables to solve for the fourth. The formula is in the reference tables.

The titration process uses a known amount and concentration of one solution to neutralize a known amount of another solution. With these three pieces of information, we can solve for the molarity of the unkown solution.

Figure 1 Before the experiment begins, the acid solution is clear. You may want to place a piece of white paper under the flask to aid in viewing the color change.	Figure 2 We must swirl the acid as we add the base to mix the solutions evenly. A small persistant pink spot indicates you are getting closer to the endpoint of the reaction.	Figure 3 When the entire solution turns the palest shade of pink, you have reached your endpoint. This pink color must persist for at least 15 seconds to be reliable.	Figure 4 If your solution turns a dark pink color then you have gone to far and have created a basic solution. You can "back titrate" too correct this or you can discard this sample.
We average the volume of base used from several trials. This average is the volume that we will use for the titration calculation.

Top

July 2, 2004