CHAPTER 12: SOLUTIONS
Part One: Introduction
A. Terminology.
1. Solution - homogeneous mixture of substances in which no settling occurs.
2. Consists of a solvent and one or more solutes.
3. Solute = is a substance being dissolved, and usually present in small amounts.
4. Solvent = medium in which the solutes are dissolved.
Solvent is usually the most abundant species in the mixture.
5. Solubility = the amount of a given solute that dissolves in a given quantity of solvent. NaCl in water, e.g., solubility of NaCl in water = 3.6 g/mL.
6. Two fluids that mix in all proportions to form a solution are said to be miscible. (alcohol and water). Oil and water would be said to be immiscible.
Figure 12.1
B. Importance.
1. Extremely important in all life – e.g., body fluids are solutions.
2. A solid, liquid, or gas can act as either solvent or solute. Usually the solvent is a liquid.
3. Sea water is an aqueous solution of many salts & some gases.
4. Tap water is an aqueous solution with trace amounts of various ions, molecular compounds, and gases. (Distilled water is pure).
Chapter 12 Page 1
5. Air is a solution of gases.
6. Solids solutions are called alloys. One example is steel, which is mostly Fe with small amounts of additives like carbon, chromium, titanium, etc. Another example is dental fillings are solid solutions called by the special name amalgams = metals dissolved in Hg.
Part Two: The Solution Process
A. Solubility and Saturated Solutions
1. Solution formation is a process of dynamic equilibrium. Example, dissolving table salt in water: NaCl(s) Na+(aq) + Cl-(aq)
Figure 12.2
Chapter 12 Page 2 Ions are continually leaving the solid crystal into the solution and returning to the crystal, until eventually:
a) all the solid is dissolved, OR b) the solution reaches its saturation point, at which some solid remains.
Which case pertains depends on the solubility of NaCl in water.
Figure 12.3
B. Factors that influence solubility.
1. General rule = “like dissolves like.”
e.g., oil dissolves in gasoline because both are hydrocarbons, which are non-polar molecules
water, however, is a polar molecule, such that oil does not dissolve in it, but alcohols do (which are also polar).
Therefore, intermolecular forces are involved.
2. There is some natural tendency of all substances to mix, based on tendency toward a state of increased disorder (ENTROPY CHANGE)
3. This is balanced by the tendency of a system to have the lowest energy possible (ENERGY CHANGE). In other words, is heat absorbed or released in forming a solution? Release of energy (exothermic) favors solution formation.
Chapter 12 Page 3 The Entropy effect and the Energy effect combine to determine the solubility. A process is favored by:
a. decrease in the energy of the system, exothermic process.
b. increase in the disorder, or randomness of the system.
4. Energy change is called the heat of solution, ΔHsolution. Depends on how strongly solute and solvent particles interact.
If ΔHsolution is negative → heat liberated (exothermic)
If ΔHsolution is positive → heat absorbed (endothermic)
The Entropy effect always works in favor of forming the solution.
5. Relative strengths of these interactions determine the extent of solubility (refer to figure below)
a. Strong solvent-solute attractions favor solubility.
b. Weak solvent-solvent attractions favor solubility.
c. Weak solute-solute attractions favor solubility.
6. Solution formation is always accompanied by an increase in the disorder of both solvent and solute, always favorable to solubility.
C. Dissolving Solids in Liquids, viewed as a stepwise process (Section 12.2)
Chapter 12 Page 4 1. Take an ionic solid. Must overcome crystal lattice energy (solute-solute interactions):
Requires large input of energy (endothermic) unfavorable to solution.
2. Must overcome solvent-solvent interactions. If solvent is water, must break Hydrogen bonds:
Requires input of energy (endothermic) unfavorable.
3. Must “solvate” the solute particles:
Exothermic, favorable to solution, especially ionic compounds dissolving in H2O.
Figure 12.8
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4. Combined Steps 2 and 3 is called “solvation,” or in case of water, “hydration.”
5. Overall energy involved: ΔHsolution =
Steps 2 & 3 Step 1
ΔHsolution = (heat of solvation) - (crystal lattice energy) OR (heat of hydration)
favorable unfavorable
Figure 12.10
6. Energy is not the only factor. Entropy increase upon solution is a favorable factor which can sometimes overcome unfavorable energy factor if latter is not too large.
+ - 7. Example: NaCl(s) + H2O(aq) → Na (aq) + Cl (aq)
ΔHsolution is (+), endothermic (unfavorable). Still occurs due to slightly larger favorable entropy consideration.
8. Ionic compounds which have very large crystal lattice energies will be insoluble in
H2O because solvation and entropy can’t overcome energy holding lattice together.
BaSO4 insoluble!
9. Some ionic solids dissolve with release of heat (exothermic) - both energy and entropy favor solution process.
CaCl2, Na2SO4, ...
Chapter 12 Page 6 10. Nonpolar solids (e.g., paraffin) do not dissolve in H2O because interactions with water (hydration) is very weak compared to water-water interactions.
Weak interaction Strong Hydrogen Weak interaction btwn. in nonpolar solid bonding solvent- solvent & nonpolar (should be easy to solvent interactions. solute. Nothing gained break up). (difficult to break up) here except entropy of HOWEVER... mixing.
Result: Water “prefers” itself rather than a non-ionic or non-polar solute. Called the “hydrophobic effect.”
D. Dissolving Liquids in Liquids – Molecular solutions (Miscibility). (Section 12.2)
1. Same factors involved. Here, though, a solid does not need to be broken up, so no lattice energy to overcome.
2. Case 1: Polar or Hydrogen-bonding solute with polar or Hydrogen bonding solvent.
Methanol dissolves in water.
H-bonding. Moderately H-bonding. Moderately H-bonding. Nothing strong solute-solute strong solvent-solvent lost, entropy of interactions. interactions. mixing favors.
3. Case 2: Nonpolar solute with polar or Hydrogen-bonding solvent.
CCl4 does not dissolve in H2O.
Weak solute-solute H-bonding. Moderately Weak solute-solvent interactions. (only strong solvent-solvent. interactions. (loss of London forces) solvent-solvent int.) (hydrophobic effect)
Chapter 12 Page 7 4. Case 3: Nonpolar solute and nonpolar solvent.
CCl4 does dissolve in benzene.
Weak solute-solute Weak solvent-solvent Weak solute-solvent interactions. (only interactions. (only (London). Entropy London forces) London forces) of mixing favors this.
5. Generalization - “like dissolves like.”
E. Gases in Liquids. (Section 14.4)
1. Same factors as before.
2. Polar gases dissolve in polar liquids, nonpolar gases dissolve in nonpolar liquids.
3. CO2 and O2 are nonpolar, so only slightly dissolve in H2O.
4. CO2 more so because of ionization:
CO2(g) + H2O(l) H2CO3(aq) carbonic acid
+ - H2CO3 (aq) H (aq) + HCO3 (aq)
5. HCl and NH3 are polar covalent gases, dissolve strongly in H2O.
Part Three: Effects of Environmental Variables on Solubility (Section 12.3)
A. Temperature Effect. (Section 12.3)
1. Explained by LeChatelier’s Principle = when a stress is applied to a system in equilibrium, the system responds in a way that relieves the stress and restores a new equilibrium.
2. When heat of solution is exothermic:
Solute + Solvent → Solution + Heat
T↑ causes shift back to the left, solubility↓.
Chapter 12 Page 8 3. When heat of solution is endothermic:
Solute + Solvent + Heat → Solution
T↑ causes shift to the right, solubility↑.
4. NaCl solubility↑ as T↑; Na2SO4 solubility↓ as T↑.
Figure 12.12
5. O2 dissolves in H2O exothermically:
T↑ O2 solubility↓ .
Basis of “thermal pollution.”
B. Pressure Effect.
1. Henry’s Law = Increase of partial pressure of a gas over a liquid increases amount of gas dissolved by a direct proportion
2. Example: Carbonated beverage stored with pressurized CO2.
3. Henry’s Law:
S = kH P
(solubility of = kH (partial pressure of gas in liquid) gas above liquid)
kH = Henry’s Law constant
Chapter 12 Page 9 Part Four: Expressions of Concentration (Section 12.4)
A. Molarity.
1. Molarity = moles of solute Liters of solution
B. Mass Percentage of solute.
1. Mass % solute = mass of solute x 100% Mass of solution
2. Example problem: A certain solution of an alcohol and water is 2.0 M alcohol and has a density of 0.95 g/mL. The alcohol has a molar mass of 126 g/mol. What is the mass percent alcohol in this solution?
C. Molality.
1. m = moles of solute kg of solvent
2. Example: What is the molality of a solution in which 1.0 gram glucose (C6H12O6) has been added to 100 mL of H2O?
3. Note that: molality ≈ molarity
for very dilute aqueous solutions.
Why? Because under these conditions:
(Liters of solution) ≈ (kg of solvent)
D. Mole Fraction.
1. XA = moles of component A total moles of all components
Chapter 12 Page 10 2. Unitless.
3. Example: What is Xglucose and Xwater in previous problem?
Part Five: Colligative Properties of Solution (Sections 12.5 through 12.8)
A. Definition.
1. Colligative property = a physical property of a solution that depends upon the number, but not the identity, of solute particles in a given amount of solvent.
2. Examples: a. vapor pressure lowering b. boiling point elevation c. freezing point depression d. osmotic pressure
3. Depends on: total molar concentration of solute particles, not on their identity.
B. Vapor Pressure Lowering - Raoult’s Law. (Section 12.5)
1. A solution containing a nonvolatile solute has a lower vapor pressure than the pure solvent.
(“nonvolatile solute” means solute has no vapor pressure of its own)
2. Physical Picture:
3. V.P. lowering depends on fraction of surface blocked by solute particles, and thus, mole fraction of solute and solvent, hence:
Chapter 12 Page 11 4. Raoult’s Law:
Psolvent = Xsolvent P°solvent
↑ ↑ ↑ v.p. of mole v.p. of pure solvent fraction solvent in solution
5. Also could write:
ΔPsolvent = Xsolute P°solvent v.p. lowering
6. Solutions obeying this perfectly are called ideal solutions.
7. Slight deviations occur when:
-- attractions differ from -- attractions and -- attractions.
C. Boiling Point Elevation. (Section 12.6)
1. Closely related to vapor pressure lowering, because boiling occurs when
V.P. of liquid = surrounding pressure.
2. As conc. solute↑ V.P. of solution↓ boiling point↑.
3. ΔTb = Kbm
where: ΔTb is boiling point elevation Kb is boiling point elevation constant (of the solvent) m is molality of solute (must be nonvolatile)
4. Each solvent has its own characteristic Kb, independent of what the solute is. See Table 12.3.
Kb = 0.512°C/molal for H2O
5. Problem: What is the normal boiling point of a glucose solution in which 270 g
glucose are added to 1.0 L of H2O?
D. Freezing Point Depression. (Section 12.6)
Chapter 12 Page 12 1. ΔTf = Kfm where: ΔTf is freezing point lowering Kf is freezing point depression constant (of solvent) m is molality of solute
2. Kf = 1.86°C/molal for H2O.
3. What is freezing point of the preceding glucose solution?
4. This is how antifreeze works, as well as salt on icy roads.
E. Determination of Molar Mass by Colligative Properties. (Section 12.6)
1. Add known mass of solute to known mass of solvent. Measure f.p. lowering ΔTf. Calculate Molar Mass (M).
ΔTf = Kfm
where: ΔTf is measured Kf is known for your solvent m must be calculated:
grams solute M solute m = kg solvent
The only unknown variable is M solute. Solve:
€ K f • grams solute M solute = ΔTf • kg solvent
2. Problem:
When 0.154 g of€ sulfur is finely ground and melted with 4.38 g of camphor, the freezing point of the camphor is lowered by 5.47°C. What is the molecular weight of
sulfur? What is its molecular formula? Given Kf = 40°C/m
Chapter 12 Page 13 F. Colligative Properties of Ionic Solutions. (Section 12.8)
1. When solute is an electrolyte (breaks into ions in solution), colligative properties are enhanced because of more particles present. Collig. Prop. depend on total concentration of ions produced.
2. Deviations from ideal behavior can be large since ions interact with each other in solution and thus are not independent.
e.g. You don’t get quite a doubling of b.p. elevation when you dissolve 1 mole of NaCl vs. 1 mole of glucose because the Na+ and Cl- ions interact.
3. Effect of ionization is expressed by the van’t Hoff factor i:
ΔT i = actual T Δ non−elec
i = 2 if NaCl had no ion-ion interactions.
€ i = 1.83 in 1.0 m NaCl
ΔTb = i Kbm modified F.P. equation for ionic solutes
G. Osmotic pressure. (Section 12.7)
1. Osmosis = spontaneous process in which solvent molecules pass through a semipermeable membrane from a solution of lower conc. to solution of higher conc.
2. Semipermeable membrane = allows solvent to flow but not solutes. (e.g., cell membranes are semipermeable)
3. osmotic pressure π = pressure that must be applied on a given solution to keep it from exerting osmotic flow on its surroundings.
Chapter 12 Page 14 4. Working equation:
π V = nRT (note similarity to ideal gas law)
π = (n/V) RT
π = MRT, where M=molarity of solute
Part Six: Colloids (Sections 12.9)
A. Definition.
1. Remember our classification of mixtures as:
a. homogeneous - true solutions like salt water, solute particles are uniformly distributed ions or molecules.
b. heterogeneous - like sand in water, particles are large enough to settle out.
2. Colloids are somewhere between these two extremes.
3. Collodial dispersion = dispersed particles are small enough to remain suspended in the solvent, so no settling takes place, but yet are large enough to make the sample cloudy.
4. Cloudiness is due to Tyndall effect = scattering of light by suspended particles when they become large enough that they are comparable to the wavelength of the light.
B. Types of colloids - study Table 12.4.
Soap molecules solubilize oils in water by forming Colloids with the oil.
oil-soluble end (Hydrophobic) water-soluble end (Hydrophilic)
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Figure 12.30 - a micelle particle
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Figure 12.31 – how soap cleans fabrics
Figure 12.34 – cell membranes
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NOTES:
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