© 2008 Brooks/Cole 1 Chemistry: The Molecular Science Moore, Stanitski and Jurs Chapter 15: The Chemistry of Solutes and Solutions

© 2008 Brooks/Cole 1

Chemistry: The Molecular ScienceChemistry: The Molecular ScienceMoore, Stanitski and Jurs

Chapter 15: The Chemistry of Solutes and Solutions

Chapter 15: The Chemistry of Solutes and Solutions


Solution = homogeneous mixture of substances. It consists of:

solventsolvent - component in the greatest amount

(or the one that does not change phase).

solutesolute - other component(s)

solvent-solute interactions determine if a substance will dissolve in a particular solvent.

Solubility & Intermolecular Forces


Solutions• Exist in all 3 physical states• Can be mixtures of solids, liquids and gases

Type of Solution Examples

Gas in gas Air

Gas in liquid Carbonated drinks.

Gas in solid Hydrogen in Pd metal.

Liquid in liquid Motor oil, vinegar.

Solid in liquid Ocean water, sugar-water.

Solid in solid Bronze, pewter, 14K gold.

Solubility & Intermolecular Forces


• Similar… = soluble. polar dissolves polar non-polar dissolves non-polar

““Like Dissolves Like”Like Dissolves Like”If solute and solvent intermolecular forces are:

Solute-Solvent Interactions

• Dissimilar…

= insoluble.



Substances dissolve when:

solvent-solute attraction > solvent-solvent attraction

and > solute-solute attraction


MiscibleMiscible liquids dissolve in all proportions.

e.g. ethanol and water (both H-bonded polar liquids).

ImmiscibleImmiscible liquids form distinct separate phases.

e.g. gasoline (non-polar) and water (polar).

colorless CCl4

green NiCl2(aq)

colorless C7H16 after mixing and settling



Solute-Solvent InteractionsName Formula Solubility (g/100g H2O @20°C)

methanol CH3OH miscible

ethanol C2H5OH miscible

1-propanol C3H7OH miscible

1-butanol C4H9OH 7.9

1-pentanol C5H11OH 2.7

1-hexanol C6H13OH 0.6

1-octanol C8H18OH immiscible

Solubility in water decreases as alcohols grow larger because:

solute-solute attraction grows (non-polar part gets bigger)

solute-solvent attraction stays ≈ constant (still only one -OH)

Londonforcesincreasing


Alcohols have a polar –OH head attached to a

non-polar hydrocarbon tail.

The head is hydrophilic (“water loving”)

The tail is hydrophobic (“water hating”)

As the tail gets bigger, it is harder and harder to dissolve since it becomes more hydrophobic.



vitamin C


You can overdose on vitamin A(or vitamin E) but not on vitamin C.

Why?


Enthalpy of HydrationEnthalpy of Hydration (ΔHhydration)

E released as an ion becomes hydrated.

(make bonds)

Lattice energyLattice energy E required to overcome the forces holding the ions together in a crystal.

(break bonds)

Dissolving Ionic Solids in Liquids

EntropyEntropyMixing, or increasing disorder, is a favorable process E-wise.

∆Hsoln = -U + (∆Hhyd,cation + ∆Hhyd,anion)


Solubility and Equilibrium

SolubilitySolubility:: the maximum quantity of solute that dissolves in a given quantity of solvent at a particular T (saturated soln.)

SaturatedSaturated• no more solute will dissolve• undissolved solid • dynamic equilibrium.

solute (s) solute (aq)

inc. T, inc. s


UnsaturatedUnsaturated[Solute] < solubility

all solid is dissolved

SupersaturatedSupersaturated[Solute] > solubility

all solid is dissolved



SupersaturatedSupersaturated solutions have more than the equilibrium amount dissolved. (unstable)


supersaturated saturatedThere is still solute dissolved!


Sodium Acetate LABStep…..2. room temperature3. elevated temperature4. after slow cool5. after added crystal

4 3

2,5

Solubility Curve


Temperature and Solubility

Solubility of SolidsSolubility of SolidsSolubility usually increasesincreases as T increases.

(usually endothermic +∆Hsoln)

WS!

Increase in T favorsendothermic process.

solute (s) + heat solute (aq)


Temperature and Solubility

Thermal Pollution

Solubility of GasesSolubility of GasesSolubility usually decreasesdecreases as T increases. (usually exothermic -∆Hsoln)


Pressure and Dissolving Gases

Gas solubility also depends upon the P of the gas above a liquid. Solubility increases with the P.

gas* + solvent saturated solution + heat*

Henry’s LawHenry’s Law

Sg = kH Pg


Solution Concentration: Units

ExampleExampleSaline solutions (NaCl in water) are often used in medicine. What is the weight percent of NaCl in a solution of 4.6 g of NaCl in 500. g of water?

Weight percentWeight percent = Mass fraction x 100%

Mass fractionMass fraction =Mass Solute

Total Mass of Solution

weight percent = 0.0091 x 100% = 0.91 %

mass fraction =4.6 g

500. g + 4.6 g= 0.0091


Practice

Weight percentWeight percent = Mass fraction x 100%

How would you prepare 425. mL of aqueous solution containing 2.40% by mass of sodium acetate? (Assume the density is the same as water.)

10.2g


Parts per Million

mass solutemass solution

x 106Parts per million (ppm)Parts per million (ppm) =

mass solutemass solution

Weight % = Weight % =

Parts per hundred (pph)Parts per hundred (pph) = x 102

Practice: Convert 73.2 ppm to weight percent.

2.4%

= 24,000 ppm

0.024%

= 240 ppm

0.00732%


Molarity and Molality

MolarityMolarity = M = moles of soluteliters of solution

MolalityMolality = m = moles of solutekilograms of solvent

Molality is a mass-based unit. • Uses solventsolvent mass (not solution).• It is T independent (unlike molarity)


PracticeCommercial 30.0% hydrogen peroxide has density = 1.11 g/mL at 25°C. What is its molarity? 9.79 M

Sea water is 10,600 ppm Na+. Calculate the mass percent and molarity of sodium ions in sea water. The density of sea water is 1.03 g/mL. 1.06%, 0.475M

Calculate the molarity and the molality of NaCl in a 20% aqueous NaCl solution whose density is 1.148 g/mL at 25C.

4.28m, 3.93M

15.12 The “proof” of an alcoholic beverage is defined as twice the percent by volume of alcohol in the beverage. What is the molarity of ethanol, C2H5OH in 1L of 90-proof beverage? Dethanol=0.79g/mL. Dbeverage=0.861g/mL. 15.2m


Colligative Properties

Properties that depend only on the concentration of solute particles (ions or molecules) in the solution and not the type.

•VP lowering (Raoult’s Law)•Freezing depression• BP elevation•Osmosis


Vapor Pressure Lowering

Solvent vapor P drops if non-volatile solute is added.

Lower purity solvent = lower vapor P.


A result of lowering the vapor pressure…

BP: Higher T is needed to get the VP = ext. P.


Colligative Properties

Boiling Point ElevationBoiling Point Elevation

ΔTb = Kb msolute

Freezing Point DepressionFreezing Point Depression

ΔTf = Kf msolute

•the Kb, Kf values only depends on the solventsolvent

• the molality, m depends on the solutesolute

Salted water raises the boiling point.

Ethylene glycol in the radiator lowers the freezing point.


colligative propertiescolligative properties…

Colligative Properties of Electrolytes

• depend upon the number of “particles” in solution.• The type of particle is unimportant.• 1 M sugar and 1 M urea aqueous solutions have

the same effect.• 1 M NaCl will be different.

Each NaCl yields 2 particles in solution

(1 Na+ ion and 1 Cl- ion).


In aqueous solution:

1 mol sucrose → 1 mol particles

1 mol NaCl → 2 mol particles (Na+ and Cl-)1 mol CaCl2 → 3 mol (Ca2+ and 2Cl-)

Modify the formulas by replacing

msolute with isolutemsolute

i = the number of particles per formula unit (van’t Hoff factor)



Boiling Point ElevationExampleExampleAt what T will 0.100 m aqueous solutions of urea, NaCl and sucrose boil? For water Kb= 0.51°C kg mol-1

b.p. of aq. urea = b.p. of aq. sucrose

ΔTb = Kb msolute = 0.51°C kg mol-1(0.100 mol/kg)

= 0.051 °C

b.p. = 100.00°C + 0.051°C = 100.05 °C


Boiling Point Elevation

The b.p. of 0.100 molal NaCl:

ΔTb = Kb iisolutesolutemsolute

= 0.51°C kg mol-1(2)(0.100 mol/kg) = 0.10 °C

b.p. = 100.00°C + 0.10°C = 100.10 °C

ExampleExampleAt what T will 0.100 m aqueous solutions of urea, NaCl and sucrose boil? For water Kb= 0.51°C kg mol-1


Freezing Point LoweringCalculate the f.p. of an aqueous 30.0% ethylene glycol mixture. For water Kf = 1.86°C kg mol-1.

100.0 g of 30% mix: 30.0 g C2H2(OH)2 + 70.0 g H2O

nglycol = 30.0 g / 62.07 g mol-1 = 0.4833 mol

mglycol = (0.4833 mol / 0.070 kg) = 6.904 molal

ΔTf = Kf msolute

ΔTf = 1.86°C kg mol-1(6.904 mol/kg) = 12.8 °C

freezing point = 0.00°C – 12.8°C = -12.8°C


Arrange these aqueous solutions in increasing b.p order: 0.10 m BaCl2

0.12 m K2SO4

0.12 m KBr

Similar molalities, so similar deviations from iexpected

m iexpected: BaCl2 = 0.10(3) = 0.30

K2SO4 = 0.12(3) = 0.36

KBr = 0.12(2) = 0.24

b.p order: KBr < BaCl2 < K2SO4




You want to purchase a salt to melt snow and ice on your sidewalk. Which one of the fjollowing salts would best accomplish your task using the least amount: KCl or CaCl2?Kf = 1.86°C kg mol-1


Practice – Calculating MM

Mass of solute = 0.00239 gMass of solvent = 0.1030 gFreezing point of a solution = 25.70°CFreeaing point of solvent 0= 26.84°CKf = 8.00 °C/m

Find the MM of the solute.


Osmotic Pressure of Solutions

Semipermeable membraneSemipermeable membrane

A thin material that• Allows passage of

small “particles”.• Stops large “particles”

OsmosisOsmosis

Movement of solventsolvent through a semipermeable membrane from dilute to more concentrated solution.

Ion surrounded by a coordination sphere of water – too large to pass through

Large molecule cannot pass.

solvent flowe.g. animal bladders and cell membranes.


Osmotic pressureOsmotic pressureOsmotic pressure = P that must be applied to stopstop osmosis.

= c R T i

Pure water

semipermeable bag of 5%

sugar water

Water enters the bag,

increasing the P…

height of the column of

solution is a measure of

c R Tn R TV

P = =


Osmotic Pressure

A cell can be exposed to 3 kinds of solution:

IsotonicIsotonic

[solute]out= [solute]in

No net flow.

HypotonicHypotonic

[solute]out< [solute]in

Net flow in.

HypertonicHypertonic

[solute]out > [solute]in

Net flow out.

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© 2008 Brooks/Cole 1 Chemistry: The Molecular Science Moore, Stanitski and Jurs Chapter 15: The Chemistry of Solutes and Solutions