chapter 12 solutions - ohio northern universitys-bates/chem172/ch12presstudent.pdfchapter 12...
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Chapter 12
Solutions
◆ solute + solvent ! solutionspeciation?stoichiometry?
◆ empirical solubility rules: Which ionic compounds are soluble in water?
◆ electrolyte definitions:strong electrolytes - soluble ionic compounds
strong acids & strong bases
weak electrolytes - weak acids & weak bases
nonelectrolytes - molecular compounds insoluble ionic compounds
Aqueous SolutionsWhat did we learn in Chapter 4?
◆ representation of how much solute is present
qualitatively: concentrated vs. dilute
quantitatively: molarity or molar concentration (units mol/L)
M = –––––––––––––––
Solution Composition
mol solutevolume solution
Types of Solutions:Physical State of Solute and Solvent
Gas Phase Solution Formation
◆ speciation
molecular compounds dissolve to produce solutions containing dispersed, individual, discrete, neutral molecules
◆ soluble molecular compounds are nonelectrolytes
Solution Formation with a Molecular Solute
◆ speciation
ionic compounds dissolve to produce solutions containing dispersed cations and anions
◆ must consider the compound’s stoichiometry to fully understand the speciation
◆ soluble ionic compounds are strong electrolytes
Solution Formation with an Ionic Solute Ionic Compounds - Soluble or Insoluble?
◆ lattice energy energy required to separate an ionic solid into its constituent gas phase ions
NaCl (s) ! Na+ (g) + Cl- (g)
◆ lattice energy increases with increasing ionic charge and decreasing ionic radius
◆ in general - the greater the lattice energy, the lower the solubility
Saturated vs. Unsaturated Solutions
overall: ∆Hsoln = ∆H1 + ∆H2 + ∆H3
◆ when solutions form, there is an increase in disorder increase in entropy
◆ solution formation is a spontaneous process
◆ sol’n formation may be endothermic or exothermic
Thermodynamics of Solution Formation
◆ representation of how much solute is present
qualitatively: concentrated vs. dilute
quantitatively: molarity or molar concentration (units mol/L)
M = –––––––––––––––
Solution Composition
mol solutevolume solution
Solution Composition
◆ mass percent composition:
mass % component = ––––––––––––– x 100
◆ related to mass percent composition:
parts per million (ppm) = –––––––––––––– x 106
parts per billion (ppb) = –––––––––––––– x 109
mass componentmass of solution
mass componentmass of solution
mass of solutionmass component
Solution Composition
◆ mole fraction composition, ! (unitless):
! of component = ––––––––––––––––––––
◆ if solution is composed of 2 substances A and B:
!A + !B = 1
mol componenttotal mol of all components
Solution Composition
◆ molality or molal concentration (units mol/kg):
molality, m = –––––––––––
◆ calculations that involve relating a solution’s molar concentration and molal concentration will require information about the density of the solution
mol solutekg solvent
example:
A solution is prepared by dissolving 13.5 g glucose (C6H12O6) in 0.100 kg of water. Calculate the mass % glucose in this solution, and the molality of the solution.
example:
A 2.5 g sample of groundwater contains 5.4 μg of Zn2+. Calculate the ppm of Zn2+ is the sample.
example:
What is the molal concentration of a solution prepared by dissolving 36.5 g C10H8 in 425 g C7H8?What is the mole fraction of C10H8 (!C10H8) in this solution?
example:
A solution of HCl (aq) is 36% HCl by mass, and has a density of 1.034 g/mL. Determine the molality and molarity of this solution.
Speciation and Solution CompositionFactors Affecting Solubility
◆ nature and strength of solute & solvent interactions
like dissolves like
◆ temperature
◆ pressure
Solute-Solvent Interactions
◆ relative strengths of attractive forces between solute and solvent particles determines if a solution will form
◆ polar solute + polar solventdipole-dipole forcepotential for H-bonding forces
◆ nonpolar solute + nonpolar solvent London dispersion forces
◆ polar solutes dissolve best in polar solvents nonpolar solutes dissolve best in nonpolar solvents
like dissolves like
Miscible and Immiscible Liquids◆ liquids that are miscible mix together in all proportions;
infinite solubility◆ liquids that are immiscible do not dissolve in one another;
zero solubility
Hydrophilic vs. Hydrophobic Compounds
◆ water-soluble compounds:hydrophilic (water loving)more polar functionality
◆ water-insoluble compounds:hydrophobic (water fearing)more hydrocarbon (nonpolar) functionalityhydrophobic compounds are lipophilic (fat souble)
vitamin C
vitamin A
example:
Predict whether each of the following will be more soluble in CCl4 or C2H5OH:
C7H16 Na2SO4 HCl I2
Effect of Temperature on Solubility
◆ the solubility of most solid solutes increases as temperature increases
◆ the solubility of gases decreases as temperature increases
Effect of Pressure on Solubility
◆ the solubility of a gas increases with increasing pressure
◆ relationship between P and solubility is given by Henry’s Law:
C = kP
C = concentration of dissolved gas
k = Henry’s Law constantP = partial pressure of gas
over solution
example:
Calculate the concentration of CO2 (g) in a carbonated drink bottled under PCO2 = 4.0 atm at 25°C. For CO2 (g) in water at 25°C, k = 0.031 mol/L•atm.
Then - calculate the concentration of CO2 (g) in the drink after opening if PCO2 decreases to 3.0 x 10–4 atm.
Properties of Solutions
◆ How do the following properties change when a solute is dissolved in a solvent?
vapor pressurefreezing pointboiling pointosmotic pressure
◆ We will be discussing colligative properties - dependent on the collection of particles present.
◆ We will need to consider the nature of solutes used: volatile or nonvolatile? electrolyte or nonelectrolyte?
◆ The addition of a nonvolatile solute to a solvent results in a solution with a lower vapor pressure (Psol’n) than the vapor pressure of the pure solvent (P°solv).
◆ Raoult’s Law:
Psol’n = (!solv)(P°solv)
◆ as amount of solute increases:mol solute increases!solute increases!solvent decreasesPsol’n decreases
Vapor Pressure of Solutions:Raoult’s Law
Ideal Solutions vs. Nonideal Solutions
positive deviations◆ weaker solute-
solvent interactions◆ see higher than
predicted Psol’n
negative deviations ◆ stronger solute-
solvent interactions◆ see lower than predicted Psol’n
◆ an ideal solution follows Raoult’s Law
◆ nonideal solutions deviate from Raoult’s Law behavior
example:
Calculate the vapor pressure of a solution prepared by dissolving 50.0 mL of glycerin in 500.0 mL water at 25°C.
some details:◆ glycerin, C3H8O3 is a nonvolatile solute
molar mass = 92.09 g/mol d = 1.26 g/mL
◆ for water at 25°C, P°vap = 23.8 torr molar mass = 18.02 g/mol d = 1.00 g/mL
Solution with 2 Volatile Components
◆ each component will follow Raoult’s law:PA = !AP°A and PB = !BP°B
Ptot = PA + PB = !AP°A + !BP°B
example:
Consider a mixture of 1.0 mol C6H6 + 2.0 mol C7H8.Determine the vapor pressure of this solution. For C6H6, P° = 75 torr; for C7H8, P° = 22 torr.
Boiling Point Elevation & Freezing Point Depression
◆ the boiling point of a solution is higher than the boiling point of the pure solvent
∆Tb = Kb•m
◆ the freezing point of a solution is lower than the freezing point of the pure solvent
∆Tf = Kf•m
Boiling Point Elevation & Freezing Point DepressionSet-up for Freezing Point Depression Experiment:
∆Tf = Kf•m
example:
A solution is prepared by dissolving 42.0 g eucalyptol (C10H18O, molar mass = 154.2 g/mol) in 0.600 kg chloroform (CHCl3).
Determine the boiling point and freezing point of this solution. For CHCl3, Kb = 3.63 °C/m; bp = 61.2°C; Kf = 4.68 °C/m; fp = –63.5°C .
Osmotic Pressure◆ Osmosis is the net movement of solvent molecules through
a semi-permeable membrane from a region of lower concentration to higher concentration.
◆ results in ∆volume
◆ in turn there is a change in P exerted on membrane
◆ osmotic pressure - pressure required to stop osmosis
Osmotic Pressure, ∏
∏V = nRT OR ∏ = (n/V)RT
∏ = MRT
some terminology:
◆ isotonic solutions – solutions with the same osmotic pressure (and, ∴ same M at constant T)
◆ hypotonic solution – solution with a lower ∏ and M
◆ hypertonic solution – solution with a higher ∏ and M
example:
The average osmotic pressure of blood at 25°C is 7.70 atm. Determine the molar concentration of C6H12O6 (aq) that is isotonic with blood.
What is the molar concentration of NaCl (aq) that is isotonic with blood at 25°C?
Using Solution Properties forDetermination of Molar Mass of a Solute
example:
0.250 g of a nonvolatile nonelectrolyte solute is dissolved in 40.0 g CCl4. The boiling point for the resulting solution is determined by experiment to be 77.157°C.
Determine the molar mass of the solute.
for CCl4: bp = 76.800°C; Kb = 5.02 °C/m
Using Solution Properties forDetermination of Molar Mass of a Solute
example:
3.50 mg of a protein is dissolved in enough water to produce 5.00 mL of solution. The measured osmotic pressure of the solution is 1.54 torr at 25°C.
Determine the molar mass of the protein.
Properties of Electrolyte Solutions
◆ properties of solutions depend on the total concentration of solute particles:
solutionparticle
concentration expected fp observed fp
0.100 m C6H12O6 (aq) 0.100 m – 0.186°C – 0.186°C
0.100 m NaNO3 (aq)
0.200 m – 0.372°C –0.348°C
0.100 m K2SO4 (aq) 0.300 m – 0.558°C –0.430°C
◆ note: observed fp of electrolyte solutions is less than the expected (calculated) fp
Properties of Electrolyte Solutions
◆ the observed difference in properties of electrolyte solutions is attributed to ion pairing
◆ total ion concentration is actually slightly lower than predicted based on stoichiometric considerations of complete dissociation of compound
◆ van’t Hoff factor, i – gives a measure of the extent of electrolyte dissociation
i = ––––––––––mol particlesmol solute
Properties of Electrolyte Solutions
solutetheoretical value of i
observed i0.10 m
observed i0.010 m
observed i0.0010 m
C6H12O6 1.00 1.00 1.00 1.00
NaCl 2.00 1.87 1.94 1.97
MgSO4 2.00 1.21 1.53 1.82
K2SO4 3.00 2.31 2.70 2.84
note:◆ i closer to theoretical value at lower concentrations;
extent of ion pairing is less in more dilute solutions
◆ i closer to ideal value when ion charges are smaller extent of ion pairing is less between ions of lower charge