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