Chem. 31 – 2/25 Lecture
Announcements I
• Exam 1
– On Monday (3/2)
– Will Cover the parts we have covered in Ch. 1, 3 and
4 plus parts of Ch. 6 (through Le Châtelier’s Principle)
– Some of HW1.3 postponed (see posted solutions)
– Help session (11:00 to 12:00 on Friday – after office
hours)
Announcements II
• Today’s Lecture
– Chapter 6 Material – Le Châtelier’s Principle (last part
on Exam1)
– Review of Material on Exam 1 (including Equation)
– Chapter 6 Material not on Exam 1 (Sparingly soluble
salts)
Le Châtelier’s Principle
Stess Number Two: Dilution
Side with more moles is favored at
lower concentrations
Example: HNO2(aq) ↔ H+ + NO2If solution is diluted, reaction goes
to products
If diluted to 2X the volume:
K
H NO 



2
HNO2 
  
1  1
H
NO2
2
Q 2
1
HNO2 
2

1
Q K
2
So Q<K, products favored
Le Châtelier’s Principle
Stess Number Two: Dilution – Molecular Scale View
Concentrated Solution
H+ NO2-
H+ NO2-
Diluted Solution – dissociation allows
ions to fill more space
H+ NO2-
H+ NO2-
H+
H+
H+ NO2H+ NO2NO2-
H+ NO2-
H+ NO2NO2-
Le Châtelier’s Principle
Stress Number 3: Temperature
If ΔH>0, as T increases, products favored
If ΔH<0, as T increases, reactants favored
Easiest to remember by considering heat a
reactant or product
Example:
OH- + H+ ↔ H2O(l) + heat
Increase in T
Some Le Chatelier’s Principle Examples
•
Looking at the reaction below, that is initially
at equilibrium,
AgCl(s) ↔ Ag+(aq) + Cl-(aq) (ΔH°>0)
determine the direction (toward products or
reactants) each of the following changes will
result in
a)
b)
c)
d)
increasing the temperature
addition of water
addition of AgCl(s)
addition of NaCl
Review for Exam
•
Know the following (from Ch. 1)
–
–
–
–
–
–
Common base units (m, kg, s, mol, K) + common
multipliers (nano to kilo)
How to convert between different units*
Definitions of main concentration units (M, weight
fractions including % and ppm, and mass/volume)
How to convert between concentration units*
Steps to make solutions of known concentration +
calculations for solution preparation*
How to do stoichiometry problems (involving solids
or solutions)*
Note: *means need quantitative knowledge
Review for Exam – cont.
• Know the following (from Ch. 1 – cont.)
– Titration definitions (titrant, equivalence point, end
point, standardization titrations, analyte titrations,
back titrations)
– Practical titration requirements
– How to solve normal and back titration problems*
• Know the following (from Ch. 3)
– Rules for significant figures (including for calculations
with +, -, *, or / and when uncertainties are given)*
– Definitions for: systematic and random error,
accuracy and precision, uncertainty, relative error and
relative uncertainty
Review for Exam
• Know the following (from Ch. 3 – cont.)
-How to do propagation of uncertainty problems (+, -,
*, /, exponent, and mixed operations) and to convert
between absolute and relative uncertainty*
•Know the following (from Ch. 4)
-What a Gaussian distribution represents
-How to calculate mean values and standard deviations
(can use calculators)*
-The differences between populations and samples
-How to calculate Z values*
Review for Exam – Ch. 4 (cont.)
•How to use Table 4-1 and Z values to calculate
probabilities between limits*
•How to determine confidence intervals* + factors which
influence confidence intervals
•Difference between Z and t based confidence intervals
(lecture only)
•What confidence intervals tell you
•How to perform t-test (case 1)*
•How to recognize and select a proper test (3 t tests, F
test and Grubbs test)
•How to deal with poor data points (including use of
Grubbs test)*
Review for Exam
• Chapter 4 – cont.
-
How method of least squares works (qualitatively)
Steps to the calibration process
Assumptions required for least squares analysis
How to determine concentrations of unknowns* +
limitations in this
• Chapter 6
- Be able to write equilibrium equations from given
equilibrium reactions
- Manipulate equilibrium reactions/equations*
- Definitions of changes in Enthalpy, Entropy and Free
Energy plus predictions given reaction
Review for Exam – Ch. 6 (cont.)
•Chapter 6
•Be able to determine K from DG° or visa versa
•Be able to calculate DG from DH and DS
•Be able to predicts shift in equilibrium due to changes
in conditions (Le Châtelier’s Principle)
Review for Exam
• Equations I will give
- Basic propagation of uncertainty equations (e.g. for Y
= a + b, Y = a·b, and Y = xn)
- Equation for calculation of standard deviation
- Grubb’s test equation
- Equation for converting K to DGº
Ch. 6 – Solubility Problems
•Why Solubility is Important
• Use in gravimetric analysis (predict if precipitation is
complete enough)
• Use in precipitation titrations (not covered)
• Use in separations (e.g. separation of Mg2+ from Ca2+ in tap
water for separate analysis)
• Understand phase in which analytes will exist
•Problem Overview
• Dissolution of sparingly soluble salts in water
• Dissolution of sparingly soluble salts in common ion
• Precipitation problems (and selective precipitation problems)
Solubility Product Problems
- Solubility in Water
Example: solubility of Mg(OH)2 in water
Solubility defined as mol Mg(OH)2 dissolved/L sol’n
or g Mg(OH)2 dissolved/L sol’n or other units
Use ICE approach:
Mg(OH)2(s) ↔ Mg2+ + 2OHInitial
0
0
Change
+x
+2x
Equilibrium
x
2x
Note: x = [Mg2+] = solubility
Solubility Product Problems
- Solubility of Mg(OH)2 in water
Equilibrium Equation: Ksp = [Mg2+][OH-]2
Ksp = 7.1 x 10-12 = x(2x)2 = 4x3 (see
Appendix F for Ksp)
x = (7.1 x 10-12/4)1/3 = 1.2 x 10-4 M
Solubility = 1.2 x 10-4 M = [Mg2+]
Conc. [OH-] = 2x = 2.4 x 10-4 M
Solubility Product Problems
- Solubility of Mg(OH)2 in Common Ion
If we dissolve Mg(OH)2 in a common ion
(OH- or Mg2+), from Le Châtelier’s
principle, we know the solubility will be
reduced
Example 1) What is the solubility of
Mg(OH)2 in a pH = 11.0 buffer?
No ICE table needed because, from pH, we
know [OH-]eq and buffer means dissolution
of Mg(OH)2 doesn’t affect pH.
Solubility Product Problems
- Solubility of Mg(OH)2 at pH 11 – cont.
[H+] = 10-pH = 10-11 M
and [OH-] = Kw/[H+] = 10-3 M
Ksp = [Mg2+][OH-]2
Moles Mg(OH)2 dissolved = moles Mg2+
[Mg2+] = Ksp/[OH-]2 = 7.1 x 10-12/(10-3)2
[Mg2+] = 7 x 10-6 M
Solubility Product Problems
- Solubility of Mg(OH)2 in Common Ion
Example 2) Solubility of Mg(OH)2 in 5.0 x
10-3 M MgCl2.
Solubility Product Problems
Precipitation Problems
What occurs if we mix 50 mL of 0.020 M
BaCl2 with 50 mL of 3.0 x 10-4 M
(NH4)2SO4?
Does any solid form from the mixing of
ions?
What are the concentrations of ions
remaining?