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ChemActivity #5
Acids and Bases
Model 1: Two Definitions of Acids and Bases.
Arrhenius Definitions:
An acid is a substance that produces hydronium ions, H3O+(aq), when it is added to
water.
A base is a substance that produces hydroxide ions, OH-(aq), when it is added to
water.
Bronsted-Lowry Definitions:
An acid is a substance that donates a proton, H+, to another species.
A base is a substance that accepts a proton, H+, from another species.
Acid-base reactions are one of the most important types of chemical reactions.
Table 1. Some examples of acid-base reactions.
Reaction 1: HCl (g) + H2O (l)) ( H3O+(aq) + Cl- (aq) (Ka = 1.0 X 10+7)
Reaction 2: NH3 (aq) + H2O (l)) ( NH4+(aq)+ OH- (aq) (Kb = 1.8 X 10-5)
Reaction 3: HCN (aq) + H2O (l)) ( H3O+ (aq)+ CN- (aq) (Ka = 6.3 X 10-10)
Determining the strength of an acid or base
The strength of an acid or base can be assessed by evaluating the equilibrium position of the reaction. Reactions with an equilibrium position to the right towards products are strong, while those whose equilibrium positions that lie to the left towards reactants are weak. We can quantitate the equilibrium position of a reaction by using the equilibrium constant K, where
K = [products]/[reactants].
For acids, we use a Ka: For bases, we use a Kb:
HA + H20 ( H30+ + A- B + H20 ( BH+ + OH-
Ka = [H+][A-]/[HA] Kb = [BH+][OH-]/[B]
We can determine if an acid (or base) is strong or weak by looking at the exponent of the Ka (or Kb). If the exponent is >1, the top half of the equation dominates so the equilibrium position lies towards products indicating a strong acid (or base). Likewise, if the exponent is <1, the bottom half of the equation dominates so the equilibrium position lies towards reactants indicating a weak acid (or base).
Critical Thinking Questions:
1. a) What chemical species are the Arrhenius acids in the forward reactions (1)-(3)?
b) What chemical species are the Arrhenius bases in the forward reactions (1)-(3)?
c) What chemical species are the Bronsted-Lowry acids in the forward reactions
(1)-(3)?
d) What chemical species are the Bronsted-Lowry bases in the forward reactions
(1)-(3)?
2. Is it possible for a substance to act as both an acid and a base? Explain your reasoning.
3. Which do you think is considered the stronger acid, HCl or HCN? Explain your reasoning.
4. Consider reaction (1).
a) What species results from the loss of a proton from the Bronsted-Lowry acid in the forward reaction?
b) Does the species indicated in part a) (the answer that you gave) act as an acid or a base when the reverse of reaction (1) occurs?
c) What species results from the gain of a proton by the Bronsted-Lowry base in the forward reaction?
d) Does the species indicated in part c) act as an acid or a base when the reverse of reaction (1) occurs?
e) Do these same relationships exist for reactions (2) and (3)?
Model 2: Conjugate Pairs
Within the Bronsted-Lowry model, certain pairs of molecules are described as a conjugate acid-base pair. The two species in a conjugate acid-base pair differ by a proton only. A base is said to have a conjugate acid, and an acid is said to have a conjugate base.
A conjugate acid-base pair differs by a proton, H+. The species with the most protons is the acid.
Table 2. Examples of conjugate acid-base pairs.
Acid Base
H2CO3 HCO3-
HCO3- CO3-2
H3O+ H2O
H2S HS-
Critical Thinking Questions:
5. Why is the charge on the hydrogen sulfide ion in Table 2 given as -1?
6. Answer and explain each of the following:
a) What is the conjugate acid of NH3?
b) What is the conjugate base of H2O?
c) Define a conjugate acid-base pair.
Exercises:
1. Give the conjugate base of each of the following:
HS-; HCO3-; H2O; OH-; H3O+; NH4+; CH3NH3+; HF; CH3COOH
2. Give the conjugate acid of each of the following:
S-2; CO3-2; H2O; OH-; O-2; NH3; CH3NH2; CH3COO-; F-; HCO3-; NH2-
3. For each of the following five reactions:
i) H2SO4 (aq) + H2O ( H3O+ (aq) + HSO4- (aq)
ii) HSO4- (aq) + H2O ( SO4-2 (aq) + H3O+ (aq)
iii) H2O + H2O ( H3O+ (aq) + OH- (aq)
iv) HCN (aq) + CO3-2 (aq) ( HCO3- (aq) + CN- (aq)
v) H2S (g) + NH3 (1) ( HS- (am) + NH4+ (am) [(am) = dissolved in liquid ammonia]
a) Which reactant is the acid?
b) Which reactant is the base?
c) Find the two conjugate pairs present in the reaction.
4. Ammonia (NH3) can react as an acid or a base.
a) What is the conjugate acid of ammonia?
b) What is the conjugate base of ammonia?
c) Write an acid-base reaction in which NH3 (l) acts as both an acid and a base:
5. Complete the following table of conjugate acids and bases:
|
Acid |
Base |
|
H2S |
|
|
|
S2- |
|
H3PO4 |
|
|
H3O+ |
|
|
OH- |
|
|
|
NO2- |
|
HOCl |
|
|
|
OCN- |
|
|
F- |
|
|
HPO4-2 |
The water-dissociation equilibrium constant, Kw, is (at 25°C)
Kw = [H3O+] [OH-] = 1.0 x 10-14
Several definitions have been found to be useful:
pH = -log [H3O+] and [H3O+] = 10(-pH)
pOH = -log [OH-] and [OH-] = 10(-pOH)
pKw = -log Kw
In general, pX = -log X
For pX expressions involving concentrations, the concentration units are always mole/liter, but they are omitted in the calculation. Thus, for example, values for pH are unitless.
Note on significant figures. For logarithms, the number of significant figures is determined by the number of digits to the right of the decimal point. The value to the left of the decimal point indicates only the power of ten by which the number is to be multiplied. For example, if pH = 2.15, then the corresponding value for [H30+] has only 2 significant figures (7.1 x 10-3). The "2" in the value "2.1511” is not considered a significant figure in the conversion to concentration.
Model: pH Is Defined as - log[H3O+].
Critical Thinking Questions:
1. Show that the pH reading in the model is correct (given that the hydronium ion concentration is 5.0 x 10-4 M).
2. Consider a neutral aqueous solution:
a) What is the pH of a neutral aqueous solution?
b) What is the pOH of a neutral aqueous solution?
3. What values of pH characterize:
a) an acidic solution?
b) a basic solution?
4. What is the numerical value of pKw?
5. Recall that log (A x B) = log A + log B. What is the relationship between pH. pOH. and pKw?
Exercises:
1. Determine the pH of each of the following solutions, and indicate whether each is acidic, basic, or neutral:
a) 0.0001 M HBr
b) 0.01 mM HNO3
c) 0.001 M NaOH
d) 1 X 10-12 M HCl
2. Determine the pH of each of the following solutions, and indicate whether each is acidic, basic, or neutral:
a) Milk, [H3O+] = 3.2 x 10-7 M
b) Pickle juice, [H3O+] = 2.0 x 10-4 M
c) Beer, [H3O+] = 3.2 x 10-5 M
d) Blood, [H3O+] = 4.0 x 10-8 M
3. Determine the [H3O+] for the acids and [OH-] for the bases of each of the following solutions, and indicate whether each is acidic, basic, or neutral:
a) Lime juice, pH = 1.9
b) Tomato juice, pH = 4.2
c) Saliva, pH = 7.0
d) Kitchen cleanser, pH = 9.3
4. Rank the following aqueous solutions in order of increasing pH without referring to a table of acid-base constants. Explain your reasoning.
pure H2O
1.0 M HCl
1.0 M NaOH
1.0 M CH3COOH
1
2
