Equilibria Electrolytes Acids and Bases review The Equilibrium Constant Equilibrium Expressions Special Equilibrium Expressions Solubility Products CommonIon Effects Weak Acids and Bases ID: 613426
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Slide1
Aqueous Equilibria
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 LawSlide2
Electrolytes……
Strong electrolytes dissociate completely in aqueous solution
NaCl, KBr, Mg(NO
3
)
2
Strong Acids, Strong Bases
Weak electrolytes dissociate or only partially react in aqueous solution
Most of these, for our examples, are weak acids or bases
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 aboveSlide3
Acids and Bases (review)
Brønsted-Lowry Definition:
Acids are proton donors
Bases are proton
acceptors
Lewis Definition
Acids are electron-pair acceptors
Bases are electron-pair
donors
Arrhenius Definition
Acids react in water to release a proton
Bases react in water to release hydroxide ionSlide4
Acid + Base
Salt + Water
Acid
+
Base
Conjugate Base
+
Conjugate Acid
Some solvents are amphiprotic
Water can act as an acid and a base!
Methanol can act as an acid and a base!
Autoprotolysis
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.Slide5
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!Slide6
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
Slide7
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 multipliedSlide8
Special Equilibrium Constants and Expressions
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)Slide9
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?Slide10
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)Slide11
“I-C-E
”
Table
Initial Conc.
(Molarity)
1
0
0
Change
(Molarity)
1
+x
+x
Equilibrium Conc.
(Molarity)
1
0 + x
0 + xSlide12
What is the ppb concentration sulfur in a saturated solution of of copper (I) thiocyanate?
Consider using an I-C-E
“
table
”
!
Write the reaction
Write the Ksp expression
Look up Ksp in standardized tableSubstitute in for ion concentrations?Solve algebraically!Concentrations are in molar (M) units, you may need to convert to ppm, ppb, etc.Slide13
What is the concentration of the salt, and each ion, in a saturated solution of Calcium Phosphate?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.Slide14
Solubility Rules (General)Slide15
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.Slide16
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.Slide17Slide18
Calculations……..
What is the pH of a 1.0 M solution of acetic acid (HAc)?
What assumption can you make?
If [acid] is about 1000 times the Ka value, it
’
s concentration in solution won
’
t change much!
Use an “I-C-E” table to look at this.There are more elaborate discussions of approximations.Slide19
What is the pH of a 4.0 M solution of phosphate ion?Write reaction
Calculate Kb
Setup
“
I-C-E
”
table
Make assumptions
Solve algebraically.Slide20
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.Slide21
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
”Slide22
Henderson-Hasselbalch (Buffer) Equation
A modification of the equation for the dissociation of a weak acid.
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?Slide23
H-H Equation & Buffers….
If [A-] = [HA] pH = pKa!
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 casesSlide24
You want 1L of a buffer system that has a pH of 3.90?What acid/conjugate base pair would you use?
How would you go about figuring out how much of each reagent you might need?
How would you prepare and adjust the pH of this solution?Slide25
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:
Slide26Slide27
What is the concentration of carbon dioxide in otherwise pure freshwater at the current partial pressure of CO
2
in the atmosphere?
Partial pressure of CO
2
= 39 Pa
K
H
= 29.4 Latm/mol1 Pa = 9.9E-6 atmWhy worry about CO2 in the atmosphere in regards to water or other solutions?