Chapter 15 Worksheet 1
Bronsted-Lowry definition of acids and bases, pH scale, strong acids and bases
Arrhenius’ definition of acids and bases:
Acid: Substance that produces aqueous H+ (protons) when dissolved in water.
Hydrochloric acid: HCl(aq) → H+(aq) + Cl-(aq) Ka is large; strong acid
Acetic acid: CH3COOH(aq) = H+(aq) + CH3COO-(aq) Ka = 1.8 x 10-5; weak acid
Base: Substance that produces aqueous OH- when dissolved in water.
Sodium hydroxide: NaOH(aq) → Na+(aq) + OH-(aq) Kb is large; strong base
(Notice that all metal hydroxides are Arrhenius Bases)
Bronsted-Lowry definition of acids and bases
An acid is a proton donor and a base is a proton acceptor. An acid-base reaction involves the transfer of a proton from an acid to a base. A reaction between an acid and a base always produces another acid and another base.
Consider the reaction of a generic monoprotic acid (HA) with water:
HA(aq) + H2O(l) = H3O+(aq) + A-(aq) Keq = Ka = [H3O+]eq[A-]eq/[HA]eq
acid 1 base 1 acid 2 base 2
Ka is the equilibrium constant for the reaction of an acid with water. It is called the acid-ionization constant. The size of Ka is a measure of acid strength.
Often, H+ (hydrogen ion) is used as shorthand for H3O+ (hydronium ion) and acid ionization reactions are written as:
HA(aq) = H+(aq) + A-(aq)
Examples:
HCl (aq) + H2O (l) → H3O+ (aq) + Cl- (aq) Ka is large; strong acid
acid 1 base 1 acid 2 base 2
CH3COOH (aq) + H2O (l) = H3O+ (aq) + CH3COO- (aq) Ka = 1.8 x 10-5; weak acid
acid 1 base 1 acid 2 base 2
Consider the reaction of a generic base (A-) with water
A- (aq) + H2O (l) = HA (aq) + OH- (aq) Keq = Kb = [HA]eq[OH-]eq / [A-]eq
Kb is the equilibrium constant for the reaction of a base with water. It called the base-ionization constant. The size of Kb is a measure of base strength.
Examples:
H- (aq) + H2O (l) → H2 (aq) + OH-(aq) Kb is large, strong base
base 1 acid 1 acid 2 base 2
NH3(aq) + H2O(l) = NH4+(aq) + OH-(aq) Kb = 1.8 x 10-5; weak base
base 1 acid 1 acid 2 base 2
Notice that water can act as an acid or a base! Substances of this type are called “amphiprotic” or “amphoteric”.
Of course, any acid can react with any base. For example:
CH3COOH(aq) + NH3(aq) = NH4+(aq) + CH3COO-(aq)
acid 1 base 1 acid 2 base 2
In every acid-base reaction, the position of the equilibrium favors transfer of the proton from the stronger acid to the stronger base.
An acid and base that differ only in the presence or absence of a proton are called a conjugate acid-base pair.
1. Complete the table and circle all amphiprotic (amphoteric) substances:
Strong / HCl / ∞ / 0 / Negligible
H2SO4 / ∞ / 0
~20 / NO3- / ~5 x 10-16
Moderately strong / 1.0 / H2O / 1.0 x 10-14 / Weak
Weak / HSO4- / 1.2 x 10-2 / 8.3 x 10-13
7.5 x 10-3 / H2PO4- / 1.3 x 10-12
HF / 7.2 x 10-4 / 1.4 x 10-11
CH3COOH / 1.8 x 10-5 / 5.5 x 10-10
4.2 x 10-7 / HCO3- / 2.4 x 10-8
H2PO4- / 6.2 x 10-8 / 1.6 x 10-7
5.5 x 10-10 / NH3 / 1.8 x 10-5
HCO3- / 4.8 x 10-11 / 2.1 x 10-4
HPO42- / 3.6 x 10-13 / 2.7 x 10-2
H2O / 1.0 x 10-14 / 1.0 / Moderately strong
Negligible / 0 / O2- / ∞ / Strong
H2 / 0 / ∞
2. What do you notice about the relative strength of an acid and its conjugate base?
Types of acids and bases
Notice that an acid can be a cation, an anion, or have no charge. Only uncharged acids are named with the word “acid”.
There are two types of weak bases:
1. Anions that are conjugate bases of weak acids (CH3COO-, HPO42-, F-, etc.)
2. Uncharged molecules that contain lone pairs of electrons. Many weak bases contain nitrogen (NH3, amines). The proton binds through the lone pair on the nitrogen. Their conjugate acids are cations.
Notice that the conjugate base of a polyprotic acid is amphiprotic (amphoteric).
3. Complete the following acid base reactions, indicate the conjugate acid-base pairs, and state whether the reaction has a large or small equilibrium constant.
a. CH3COOH(aq) + NH3(aq) =
b. H2CO3 (aq) + NO3- (aq) =
c. H2CO3 (aq) + H2O (l) =
d. HCO3- (aq) + H2O (l) =
(HCO3- is amphoteric!)
e. NH3 (aq) + H2O (l) =
4. a. The equilibrium constant for which of the above reactions is an example of either a Ka (acid
ionization constant) or a Kb (base ionization constant)?
b. Write the balanced equation for the reaction whose equilibrium constant is the Ka for phosphoric acid.
c. Write the balanced equation for the reaction whose equilibrium constant is the Kb for the dihydrogen phosphate ion.
Autoionization (dissociation) of water and the pH scale:
H2O(l) + H2O(l) = H3O +(aq) + OH-(aq) [also written as: H2O(l) = H+(aq) + OH-(aq)]
Kw = [H3O+]eq[OH-]eq = 1.0 x 10-14 (at 25oC) MEMORIZE THIS!!
Kw is called the “ionization constant” or “ion-product constant” for water
For ANY solution, if you know the concentration of either the hydrogen ion or the hydroxide ion then you know the concentration of the other! They are ALWAYS related by the equation above!
pH = -log[H3O+] pOH = -log[OH-]
MEMORIZE THIS!!
Taking the negative log of both sides of the Kw expression gives:
pH + pOH = 14 MEMORIZE THIS!!
In pure water at 25oC:
[H3O+] = [OH-] = 10-7
Thus, pH = 7
If [H3O+] >[OH-], then pH < 7 and the solution is acidic
If [H3O+] <[OH-], then pH > 7 and the solution is basic
If [H3O+] =[OH-], then pH = 7 and the solution is neutral
5. Complete the table:
[H3O+] (M) / pH / pOH / [OH-] (M) / Acidic or Basic?3.2 x 10-8
4.20
5.50
7.8 x 10-2
Strong acids and bases (MEMORIZE THESE!!)
According to the Bronsted-Lowry definition, the reactions of strong acids and bases with water go to completion (very large equilibrium constants).
Some common strong acids:
Hydrochloric Acid HCl
Hydrobromic Acid HBr
Hydroiodic Acid HI
(HF is a weak acid; we’ll disuss this later)
Chloric Acid HClO3 (Your textbook does not consider this to be a strong acid, Ka = )
Perchloric Acid HClO4
Nitric Acid HNO3
Sulfuric Acid H2SO4 (Only first ionization goes to completion)
Moderately strong acid:
Hydronium ion H3O+
Some common strong bases:
Hydride ion H-
Oxide ion O2-
Moderately strong base:
Hydroxide ion OH-
Note: According to the Arrhenius definition of a base, soluble metal hydroxides (NaOH, KOH, etc.) are classified as strong bases since they are strong electrolytes. According to the Bronsted-Lowry definition, they are not bases at all! It is the resulting OH- ion that is the base and it is only a moderately strong base.
6. Consider what happens when the strong acid, nitric acid (HNO3), reacts with water.
a. Write the balanced equation for the ionization reaction. (There are two ways to write it.)
b. Write the two expressions for Ka.
c. What can we say about the size of Ka for this reaction?
d. What is the pH of a 0.00325 M solution of nitric acid?
e. What is the concentration of nitric acid if the pH is 1.5?
7. Consider what happens when Ca(OH)2, dissolves in water.
- Write the balanced equation for the reaction.
- What can we say about the size of the equilibrium constant for this reaction?
- What is the [OH-] in a 3.25 mM solution of calcium hydroxide? pOH?
- What is the pH of the above solution? [H3O]?
- What is the concentration of a calcium hydroxide solution with pH = 9.25?
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