Week 10: Lectures 28 – 30
Lecture 28: W 10/26
Lecture 29: F 10/28
Lecture 30: M 10/31 (Halloween Demo Show)
Solutions
Solutions: Homogeneous mixtures of 2 or more
substances; uniformly mixed on a molecular level
Reading:
Solute: what dissolves (a change in state [or
phase]); typically the minor component
BLB Ch 4.1, 4.5, 13.1 – 13.4
Homework:
Solvent: what the solute is dissolved in (unchanged
state); the major component
BLB 4: 3, 37, 72; Supp 4: 1 – 5;
BLB 13: 7, 21, 23; Supp 13: 1 - 12
Aqueous Solution: solvent is water
Reminder:
Angel Quiz 9 due on Thur 10/27
ALEKS Objective 10 due on Tue 11/1
Jensen Office Hour: 501 Chemistry Building
Tuesdays & Thursdays, 10:30 – 11:30 am
Exam 3: Mon, Nov. 7, 6:30 – 7:45 pm
Jensen
Chem 110 Chap 13
Page: 1
Critical aspects with regard to solutions
• how does the solute exist in the solution
free ions? = electrolyte (weak or strong)
no ions? = nonelectrolyte
• what goes into the solution
concentration
solubility (BLB Table 4.1)
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Solute: what type of compound?
The Solution Process
Metal + nonmetal (usually)
Solutions form when the solute – solvent IM forces
are comparable in magnitude and nature to the
combined solute – solute and solvent – solvent IM
forces
e.g. NaCl
1. Break solute – solute interactions (!H1)
• Ionic compounds: compounds consisting of
positive and negative ions
Mg(NO3)2
2. Break solvent – solvent interactions (!H2)
Na2SO4
" Energy is ____________
NH4Cl
3. Form solute – solvent interactions (!H3)
• Molecular compounds: compounds consisting of
individual molecules
All nonmetals or nonmetals and metalloids
" Energy is ____________
Processes occur
spontaneously when:
! energy is released
(exothermic)
e.g. H2O
HCl
! disorder increases
CH3COOH
NH3
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Enthalpy Changes for the Solution Process
Dissolution of NaCl in water
The enthalpy change of the overall process
depends on !H for each of these steps
Left: Net exothermic process (!Hsoln<0)
Example: NaCl
Ion-dipole attractions between the ions and water
molecules are sufficiently strong to pull the ions
from their positions in the crystal.
NaCl(s) + H2O # Na+(aq) + Cl—(aq)
Right: Net endothermic process (!Hsoln>0)
Example: NH4NO3
A process can be endothermic and spontaneous
when the increase in disorder is large enough.
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Chem 110 Chap 13
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Solution Process: Dynamic equilibrium
Solute
+
Solvent
dissolve
Solution
crystallize
Factors that affect solubility
! Type of intermolecular interactions and their
strengths (LIKE DISSOLVES LIKE)
! Temperature
" Effect on solubility of ionic solids
" Effect on solubility of gases
forward rate = backward rate
Solubility: maximum amount of a substance that
can be dissolved in solvent under given conditions
! Pressure
I. Effect of IM Forces (Like dissolves like)
Three different types of solutions:
Unsaturated solution contains less than
maximum concentration of solute; more solute
can dissolve if added to the solution
• Polar solvents dissolve polar & ionic solutes
•
Saturated solution contains maximum
concentration of solute under the given
conditions; In dynamic equilibrium with
undissolved solute; additional solute will NOT
dissolve if added to the solution
Molecule
•
Supersaturated solution contains more than
maximum concentration of solute; Solution is
unstable (that is NOT at equilibrium)
•
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Chem 110 Chap 13
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• Nonpolar solvents dissolve nonpolar solutes
solubility in water
at 25°C (g/100g of H2O)
CH3OH
total
CH3CH2OH
total
CH3CH2CH2OH
total
CH3CH2CH2CH2OH
8.06
CH3CH2CH2CH2CH2OH
2.82
CH3CH2CH2CH2CH2CH2OH
0.62
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Chem 110 Chap 13
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Demo
Solutes: I2
CuSO4
IM forces
LDF
ion-ion
Solvents: pentane (C5H12) LDF
water
LDF,
dipole-dipole,
H-bonding
Predictions?
Will I2 dissolve in pentane?
A.
B.
C.
D.
E.
Will CuSO4 dissolve in pentane?
Will CuSO4 dissolve in water?
Will pentane and water mix?
Chem 110 Chap 13
A.
B.
C.
D.
E.
H2O(l)
CH3OH(l)
HCl(l)
CH3CH2OH(l)
heptane (C7H16) (l)
2. Which one of the following substances is
most likely to dissolve in water?
Will I2 dissolve in water?
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Practice Examples:
1. Which one of the following will be most
soluble in benzene [C6H6 (l)]?
Page: 9
HOCH2CH2OH
CHCl3
CH3(CH2)9OCH3
CH3(CH2)8CH2OH
CCl4
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Chem 110 Chap 13
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II. Effect of Temperature on Solubility
• For ionic solids, solubility usually increases as
temperature increase
• For gases, solubility decreases as temperature
increases
III. Effect of Pressure on Solubility
Example: A 40 g sample of potassium nitrate is
dissolved in 100 g of water at 50°C, with
precautions taken to avoid water evaporation.
The solution is cooled to 20°C and no precipitate
is observed. This solution is ______.
Chem 110 Chap 13
• However, the solubility of gas (Cg, note: BLB uses
Sg) is proportional to the partial pressure of that
gas (Pg) above the solution.
A. Unsaturated
B. Saturated
C. Supersaturated
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• Pressure does not have any effect on the
solubility of solid or liquid;
Henry’s Law: Cg = k Pg
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Example: What is solubility for He in water at
30°C, if its partial pressure above the solution is
350 torr? The solubility constant for He in water
at this temperature is 3.7 x 10-4 M/atm.
Demo
Which Bulbs Light Up?
Bulb Wattage
7.5
25
40
pure H2O
NaCl(aq)
Practice Example: Which of the following changes
increase the solubility of N2 gas in water?
1.
2.
3.
4.
increasing the temperature
decreasing the temperature
increasing the partial pressure of the solute
decreasing the partial pressure of the solute
A.
B.
C.
D.
E.
2
4
1
2
2
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1M HCl(aq)
1M CH3COOH
sugar(aq)
CH3OH
only
only
and 3
and 4
and 3
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Electrolytes: A substance that conducts electricity
(dissociates into ions) when dissolved in water
• Strong electrolytes: good conductor
– completely ionized (in solution)
NaCl (s) + H2O # Na+ (aq) + Cl– (aq) + H2O
HCl (aq) + H2O # H3O+ (aq) + Cl– (aq)
• Weak electrolytes – poor conductor
– partially ionized (in solution)
NH3 (aq) + H2O ! NH4+ (aq) + OH– (aq)
Identifying an Electrolyte
A) Is the compound ionic?
Yes: All ionic compounds are strong electrolytes
Note: it could be insoluble
B) If the compound is molecular, is it an acid
(donate H+) or a base (accept H+) in solution?
No: it is a non electrolyte
c) If a compound is an acid or a base, is it one of
the seven strong acids or eight strong bases?
Yes: it is a strong electrolyte
No: it MUST be a weak electrolyte
HF (aq)+ H2O ! H3O+ (aq) + F– (aq)
• Non-electrolytes – non-conductors
– not ionized (in solution)
C6H12O6 (s) + H2O # C6H12O6 (aq)+ H2O
(glucose)
Note: You MUST memorize these strong acids and
bases; all are strong electrolytes!
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Summary of Electrolytes
These are all water soluble compounds. Are they
strong, weak or non-electrolytes?
• Strong electrolytes – ionic compounds, strong
acids, and strong bases that ionize completely in
solution.
A. HNO3
• Weak electrolytes – weak acids and weak bases
that are partially ionized in solution.
B. CuCl2
• Nonelectrolytes – molecular compounds that are
not acids or bases and do not ionize in solution
C. CH3OH
Examples of Organic Non-electrolytes:
D. CH3COOH (acetic acid)
Alcohols
Aldehydes
E. NH3
ketones
Carbohydrates (sugar)
F. Ca(OH)2
Examples of Organic Weak electrolytes:
Weak Acids (including Carboxylic Acids)
G. (NH4)2SO4
Weak Bases (including Amines)
I. Glucose (C6H12O6)
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Concentration Expressions
Molarity (M = ________)
mass percentage =
Example: What is the molarity of a solution
that contains 2.76 g NH4NO3 in 400 mL of
solution?
mole fraction = Xi =
A.
B.
C.
D.
E.
parts per million (ppm):
molarity (M) =
11.6 M
0.0862 M
8.62 x 10-5 M
1.16 x 10+4 M
8.62 x 10-3 M
molality (m) =
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Molality (m = ________ )
Conversion Between Concentration Units
Example: What is the molality of an aqueous
solution that is 27.0% H2SO4 by mass and
has a density of 1.198 g/mL?
Example: What is the molality of a solution
prepared by dissolving 75.6 g of C6H8O6, in
150 g of water?
A.
B.
C.
D.
E.
0.286 m
28.6 m
1.90 m
2.86 m
1.43 m
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Chem 110 Chap 13
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A.
B.
C.
D.
E.
3.77 m
0.849 m
3.30 m
32.3 m
0.275 m
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Practice Example: What is the molarity of an
aqueous solution that is 13.0% NaCl by mass
and has a density of 1.10 g/mL?
A.
B.
C.
D.
E.
143 M
2.45 M
2.56 M
2.23 M
1.43 x 10-2 M
Concentration of Strong Electrolytes
Example: Which solution contains the largest
concentration of chloride ions?
A. 100 mL of 1.0 M BaCl2
B. 40 mL of 1.0 M NaCl
C. 60 mL of 1.0 M FeCl3
Which solution contains the largest moles of
total ions?
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Concentration and Dilution
Dilution: making a weaker solution by adding
solvent
Before and after dilution:
mass (moles) of solute remains the same
moles of solute = M V
M = concentration (molarity) of solution
V = volume of solution (L) SO...
Example: A solution of sodium sulfide is
prepared by diluting 50.0 mL of 0.874 M
sodium sulfide solution to a total volume of
250.0 mL. What is the concentration of the
final solution?
A.
B.
C.
D.
E.
0.0875 M
4.37 M
0.525 M
0.350 M
0.175 M
Mi * Vi = mol solute = Mf * Vf
Concentration Dilution Problem Solving Steps:
1. Organize information
2. Draw pictures if applicable
3. Solve for unknows
4. Manipulate info to get required answer
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Practice Example: A stock solution is
prepared by adding 25 mL of 1.3 M NaNO3
to enough water to make 47 mL. How many
moles of NaNO3 are present in 35 mL of the
stock solution?
A.
B.
C.
D.
E.
0.024
0.032
0.044
0.048
0.061
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Practice Example: When 30 mL of an 80 mL
stock solution containing 1.1 M AlCl3 is
diluted to 51 mL, what is the concentration
of Cl- in the new solution?
A.
B.
C.
D.
E.
mol
mol
mol
mol
mol
Chem 110 Chap 13
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1.1 M
3.3 M
0.65 M
1.72 M
1.94 M
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