Aqueous Equilibria
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Electrolytes
Acids and Bases (review)
The Equilibrium Constant
Equilibrium Expressions
“Special” Equilibrium Expressions
Solubility Products
Common-Ion Effects
Weak Acids and Bases
Introduction to Buffers
Henry’s Law
Electrolytes……
• Strong electrolytes dissociate completely in
aqueous solution
o NaCl, KBr, Mg(NO3)2
o Strong Acids, Strong Bases
• Weak electrolytes dissociate or only partially
react in aqueous solution
o Most of these, for our examples, are weak acids
or bases
o Ammonia, ammonium, phosphoric acid (all 3
protons are weak), acetic acid/acetate ion, etc.
• Let’s write some example chemical reactions
for all of the above
Acids and Bases (review)
• Brønsted-Lowry Definition:
o Acids are proton donors
o Bases are proton acceptors
• Lewis Definition
o Acids are electron-pair acceptors
o Bases are electron-pair donors
• Arrhenius Definition
o Acids react in water to release a proton
o Bases react in water to release hydroxide ion
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Acid + Base  Salt + Water
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Acid + Base  Conjugate Base + Conjugate Acid
• Some solvents are amphiprotic
o Water can act as an acid and a base!
o Methanol can act as an acid and a base!
• Autoprotolysis
o Some solvents can react with themselves to
produce an acid and a base
• Water is a classical example
• Weak acids dissociate partially, weak bases
undergo partial hydrolysis. Strong acids and
bases are strong electrolytes.
The Equilibrium State
•Consider a generic reaction
The concentrations of each reagent are constant at equilibrium, even
though individual molecules are constantly reacting.
Concentrations are typically in molar (M) units, but gases can be
expressed as their partial pressure (atm) and solids and pure liquids will
have concentrations of unity (1).
Another way of saying this is that the reaction rate in one direction is
equal to the reaction rate in the reverse direction.
Recall Le Châtelier’s Principle and how changing reaction components
and conditions can alter equilibrium!
The Equilibrium Expression
•Consider a generic reaction
•Dissolved species are in molar (M) concentrations
•Gaseous species partial pressures are in atmospheres
•Pure liquids and pure solids have concentrations of 1.
•Excess solvents, which do NOT participate in the reaction, also
have concentrations of 1.
•Equilibrium constants are reaction, phase, temperature and
pressure dependent
Manipulating Equilibrium Expressions
• If you write a reaction in reverse, the new K is the
inverse of the original K
• If we add reactions, K values are multiplied
Special Equilibrium Constants and Expressions
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Kw (dissociation of water)
Ksp (solubility of salts in saturated solutions)
Ka (acid dissociation)
Kb (base hydrolysis)
x (complex ion formation)
KH (Henry’s Law)
Kw (Dissociation of Water)
• Water is amphiprotic
• Kw = 1.0E-14 at about 25 ˚C
• This is where the pH scale we commonly use
originates from!
• What is the concentration of hydronium and
hydroxide ions in neutral solution? What is the
pH? What is the pOH?
Solubility Products & Common Ion Effect
• Ksp applies to salts in equilibrium in
saturated solutions.
• The solution MUST be SATURATED!
• The [solid] cancels out as it is 1.
• You can calculate concentrations of the
salt, or the component ions.
• This applies to dilute solutions in pure water,
and ignores activity (we’ll not worry about
activity)
“I-C-E” Table
Initial Conc.
(Molarity)
1
0
0
Change
(Molarity)
1
+x
+x
Equilibrium
Conc.
(Molarity)
1
0+x
0+x
K sp = [Ag + ] * [Cl - ]
At equilibrium (we ALWAYS solve at Equilibrium)
K sp = [x] * [x]
• What is the ppb concentration sulfur in a
saturated solution of of copper (I)
thiocyanate?
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Consider using an I-C-E “table”!
Write the reaction
Write the Ksp expression
Look up Ksp in standardized table
Substitute in for ion concentrations?
Solve algebraically!
Concentrations are in molar (M) units, you
may need to convert to ppm, ppb, etc.
• What is the concentration of the salt, and each ion,
in a saturated solution of Calcium Phosphate?
o Consider using an I-C-E “table”!
o Write the reaction
o Write the Ksp expression
o Look up Ksp in standardized table
o Substitute in for ion concentrations?
o Solve algebraically!
o Concentrations are in molar (M) units, you may
need to convert to ppm, ppb, etc.
Solubility Rules (General)
Common Ion Effect…
• What if your now saturated solution
contained some ion before you added the
salt?
• The pre-existing “common ion” influences
the solubility of the salt!
• Use the previous steps, with an I-C-E table!
• What is the solubility of silver chloride in 1uM
sodium chloride? Setup “I-C-E” table.
Weak Acid & Weak Base Equilibria
• Weak acids produce weak conjugate
bases, and weak bases produce weak
conjugate acids
• Ka is a “special” equilibrium constant for the
dissociation of a weak acid (found in
standard tables)
• Kb is a “special” equilibrium constant for the
hydrolysis of a weak base.
Calculations……..
• What is the pH of a 1.0 M solution of
acetic acid (HAc)?
• What assumption can you make?
o If [acid] is about 1000 times the Ka value, it’s
concentration in solution won’t change
much!
o Use an “I-C-E” table to look at this.
o There are more elaborate discussions of
approximations.
• What is the pH of a 4.0 M
solution of phosphate ion?
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Write reaction
Calculate Kb
Setup “I-C-E” table
Make assumptions
Solve algebraically.
Buffers
• Buffers resist the change in pH because they have
acid to neutralize bases and bases to neutralize
acids.
• Made from a weak acid (HA) and the salt of its
conjugate base (A-, where the counter ion is gone
for example), or a weak base and the salt of its
conjugate acid.
Features of Buffers
• Buffers work best at maintaining pH near the
Ka of the acid component, usually about +/1 pH unit. This is their buffer capacity (see fig.
9-5)
• Buffers resist pH changes due to dilution.
• All seen when we use the “Buffer Equation”
Henderson-Hasselbalch (Buffer) Equation
• A modification of the equation for the
dissociation of a weak acid.
[ ]
æ A- ö
÷
pH = pKa + log ç
ç [ HA] ÷
è
ø
• The pH is the pH of the buffered solution, pKa
is the pKa of the weak acid.
• What is the pH of a buffer solution made from
1.0 M acetic acid and 0.9 M sodium acetate?
• You add .10 moles of sodium hydroxide to the
above solution? What is the new pH?
H-H Equation &
Buffers….
• If [A-] = [HA] pH = pKa!
o This is what we see at half-way to the
equivalence point in the titration of a weak
acid with a strong base!
• Dilution does not change the
ratio of A- to HA, and thus the pH
does not change significantly in
most cases
• You want 1L of a buffer system that
has a pH of 3.90?
o What acid/conjugate base pair would you use?
o How would you go about figuring out how much of
each reagent you might need?
o How would you prepare and adjust the pH of this
solution?
Henry’s Law
• At a given temperature (like any other equilibrium
situation), the amount of a gas that will dissolve in a
liquid is proportional to the partial pressure of that
gas over the liquid.
• A common form of Henry’s Law:
Pgas(atm)
KH =
[gas]molarity
• What is the concentration of carbon dioxide in
otherwise pure freshwater at the current partial
pressure of CO2 in the atmosphere?
o Partial pressure of CO2= 39 Pa
o KH = 29.4 Latm/mol
o 1 Pa = 9.9E-6 atm
• Why worry about CO2 in the atmosphere in regards
to water or other solutions?