Prentice Hall ©2004 CHAPTER 11 SOLUTIONS AND THEIR PROPERTIES Chapter 11Slide 1.

Prentice Hall ©2004

CHAPTER 11CHAPTER 11

SOLUTIONS AND THEIR PROPERTIES

Chapter 11 Slide 1

Prentice Hall ©2004 Chapter 11 Slide 2

• Saturated: Contains the maximum amount of solute that will dissolve in a given solvent.

• Unsaturated: Contains less solute than a solvent has the capacity to dissolve.

• Supersaturated: Contains more solute than would be present in a saturated solution.

• Crystallization: The process in which dissolved solute comes out of the solution and forms crystals.

Solution Formation01Solution Formation01

• Saturated: Contains the maximum amount of solute that will dissolve in a given solvent.

• Unsaturated: Contains less solute than a solvent has the capacity to dissolve.

• Supersaturated: Contains more solute than would be present in a saturated solution.

• Crystallization: The process in which dissolved solute comes out of the solution and forms crystals.









• Exothermic ∆Hsoln:

• The solute–solvent

interactions are stronger

than solute–solute or

solvent–solvent.

• Favorable process.



• Endothermic ∆Hsoln:

• The solute–solvent

interactions are weaker

than solute–solute or

solvent–solvent.

• Unfavorable process.



• Solubility: A measure of how much solute will

dissolve in a solvent at a specific temperature.

• Miscible: Two (or more) liquids that are completely

soluble in each other in all proportions.

• Solvation: The process in which an ion or a

molecule is surrounded by solvent molecules

arranged in a specific manner.



1. Predict the relative solubilities in the following cases:

(a) Br2 in benzene (C6H6) and in water,

(b) KCl in carbon tetrachloride and in liquid ammonia,

(c) urea (NH2)2CO in carbon disulfide and in water.

2. Is iodine (I2) more soluble in water or in carbon disulfide (CS2)?

3. Which would have the largest (most negative) hydration energy and which should have the smallest? Al3+, Mg2+, Na+


Concentration Units 01Concentration Units 01


• Concentration: The amount of solute present in a given amount of solution.

• Percent by Mass (weight percent): The ratio of the mass of a solute to the mass of a solution, multiplied by 100%. % bymassof solute =

mass of solute

mass of solution 100%

mass of solution =mass of solute +mass of solvent




• Parts per Million:

• Parts per million (ppm) =

= % mass x 104

• One ppm gives 1 gram of solute per 1,000,000 g or one mg per kg of solution. For dilute aqueous solutions this is about 1 mg per liter of solution.

610xsolutionofmassTotal

componentofMass



• A sample of 0.892 g of potassium chloride (KCl) is

dissolved in 54.6 g of water. What is the percent by

mass of KCl in this solution?

• An aqueous solution is 5.50% H2SO4. How many

moles of sulfuric acid (MM = 98.08 g/mol) are

dissolved in 250.0 g of the solution?



• Mole Fraction (X):

• Molarity (M):

• Molality (m):

moles of number Total Aof MolesAX

SOLUTION of Literssolute of Moles

Molarity

SOLVENT of Kilogramssolute of Moles

=Molality


Units of Concentration Units of Concentration

Mass Percent (mass %)Mass % = (mass of component/total mass of SOLUTION) x 100%

Parts per million, ppm = (mass of component/total mass of SOLUTION) x 106

Parts per billion, ppb = (mass of component / total mass of SOLUTION) x 109



• Molality from Mass: Calculate the molality of a

sulfuric acid solution containing 24.4 g of sulfuric

acid in 198 g of water. The molar mass of sulfuric

acid is 98.08 g.

• Molality from Molarity: Calculate the molality of a

5.86 M ethanol (C2H5OH) solution whose density is

0.927 g/ml.



• Molality from Mass %: Assuming that seawater is

a 3.50 mass % aqueous solution of NaCl, what is

the molality of seawater?

• Molarity from Molality: The density at 20°C of a

0.258 m solution of glucose in water is 1.0173

g/mL, and the molar mass of glucose is 180.2 g.

What is the molarity of the solution?


Concentration Units

08

Concentration Units

08

• Mole Fraction from Molality: An aqueous

solution is 0.258 m in glucose (MM = 180.2 g/mol).

What is the mole fraction of the glucose?

• Mass from Molality: What mass (in grams) of a

0.500 m aqueous solution of urea [(NH2)2CO, MM

= 60.1 g/mol] would you use to obtain 0.150 mole

of urea?


Effect of Temperatureon Solubility 01Effect of Temperatureon Solubility 01

• Solids:


Effect of Temperature on Solubility 02Effect of Temperature on Solubility 02

• Gases:


• Henry’s Law: • The solubility of a gas is proportional to the pressure of the gas over the solution.

c P

c = k·P

c kP

The Effect of Pressure on theSolubility of Gases 01The Effect of Pressure on theSolubility of Gases 01


Flash Animation - Click to ContinueFlash Animation - Click to ContinueFlash Animation - Click to ContinueFlash Animation - Click to Continue




• Calculate the molar concentration of O2 in water at 25°C for a

partial pressure of 0.22 atm. The Henry’s law constant for O2

is 3.5 x 10–4 mol/(L·atm).

• The solubility of CO2 in water is 3.2 x 10–2 M at 25°C and 1

atm pressure. What is the Henry’s law constant for CO2 in

mol/(L·atm)?


Colligative Propertiesof Nonvolatile Solutes 01Colligative Propertiesof Nonvolatile Solutes 01

• Colligative Properties: Depend only on the number of solute particles in solution. These affect properties of the solvent.

• There are four main colligative properties:1. Vapor pressure lowering

2. Freezing point depression

3. Boiling point elevation

4. Osmotic pressure



• When solute molecules displace solvent molecules at the surface, the vapor pressure drops since fewer gas molecules are needed to equalize the escape rate and capture rates at the liquid surface.



• Raoult’s Law: Psoln = P°solv Xsolv

• For a single solute solution, Xsolv= 1 – Xsolute , • We can obtain an expression for the change in vapor

pressure of the solvent (the vapor pressure lowering).Psoln = P°

solv – Psoln

= P°solv – Xsolv P°

solv

= P°solv – (1 – Xsolute ) P°

solv

∆P = Xsolute P°solv

Where superscript o is for pure substance.


Van’t Hoff FactorVan’t Hoff Factor

• For incompletely dissociating ionic solids

• Van’t Hoff Factor i = moles of particles in solution• moles of solute dissolved

•

Chapter 11 Slide 27



• The vapor pressure of a glucose (C6H12O6) solution

is 17.01 mm Hg at 20°C, while that of pure water is 17.25 mm Hg at the same temperature. Estimate the molality of the solution.

• How many grams of NaBr must be added to 250 g of water to lower the vapor pressure by 1.30 mm Hg at 40°C? The vapor pressure of water at 40°C is 55.3 mm Hg.


Colligative Properties of a Mixture of Two Volatile Liquids 01Colligative Properties of a Mixture of Two Volatile Liquids 01

• What happens if both components are volatile(have measurable vapor pressures)?

• The vapor pressure has a value intermediate between the vapor pressures of the two liquids.PT = PA + PB

= XAP°A + XBP°

B

= XAP°A + (1 – XA)P°

B

PT = P°B + (P°

A – P°B)XA


Boiling-Point Elevation and Freezing-Point Depression 01Boiling-Point Elevation and Freezing-Point Depression 01

• Boiling-Point Elevation (∆Tb): The boiling point of the solution (Tb) minus the boiling point of the pure solvent (T°

b):

∆Tb = Tb – T°b

∆Tb is proportional to concentration:

∆Tb = Kb mKb = molal boiling-point elevation constant.

Also for incompletely dissociating ionic solids

∆Tb = Kb m i



• Freezing-Point Depression (∆Tf): The freezing point of the pure solvent (T°

f) minus the freezing point of the solution (Tf).

∆Tf = T°f – Tf

∆Tf is proportional to concentration:

∆Tf = Kf m Kf = molal freezing-point depression constant.

∆Tb = Kb m i





• van’t Hoff Factor, i: This factor equals the number of ions produced from each molecule of a compound upon dissolving.

i = 1 for CH3OH i = 3 for CaCl2

i = 2 for NaCl i = 5 for Ca3(PO4)2

• For compounds that dissociate on dissolving, use:

∆Tb = iKb m ∆Tf = iKf m ∆P = ix2 P°1



• How many grams of ethylene glycol antifreeze,

CH2(OH)CH2(OH), must you dissolve in one liter of

water to get a freezing point of –20.0°C. The molar

mass of ethylene glycol is 62.01 g. For water, Kf =

1.86 (°C·kg)/mol. What will be the boiling point?



• What is the molality of an aqueous solution of KBr

whose freezing point is –2.95°C? Kf for water is 1.86

(°C·kg)/mol.

• What is the freezing point (in °C) of a solution

prepared by dissolving 7.40 g of K2SO4 in 110 g of

water? The value of Kf for water is 1.86 (°C·kg)/mol.


Osmosis and Osmotic Pressure 01Osmosis and Osmotic Pressure 01


• Osmosis: The selective passage of solvent molecules through a porous membrane from a dilute solution to a more concentrated one.

• Osmotic pressure (π or ∏): The pressure required to stop osmosis.

π = iMRT

R = 0.08206 (Latm)/(molK)








• Isotonic: Solutions have equal concentration of

solute, and so equal osmotic pressure.

• Hypertonic: Solution with higher concentration of

solute.

• Hypotonic: Solution with lower concentration of

solute.



• The average osmotic pressure of seawater is about

30.0 atm at 25°C. Calculate the molar

concentration of an aqueous solution of urea

[(NH2)2CO] that is isotonic with seawater.

• What is the osmotic pressure (in atm) of a 0.884 M

sucrose solution at 16°C?


Uses of Colligative Properties 01Uses of Colligative Properties 01

• Desalination:



• A 7.85 g sample of a compound with the

empirical formula C5H4 is dissolved in 301 g of

benzene. The freezing point of the solution is

1.05°C below that of pure benzene. What are

the molar mass and molecular formula of this

compound?



• A 202 ml benzene solution containing 2.47 g of an

organic polymer has an osmotic pressure of 8.63

mm Hg at 21°C. Calculate the molar mass of the

polymer.

• What is the molar mass of sucrose if a solution of

0.822 g of sucrose in 300.0 mL of water has an

osmotic pressure of 149 mm Hg at 298 K?



• Two miscible liquids, A and B, have vapor

pressures of 250 mm Hg and 450 mm Hg,

respectively. They were mixed in equal molar

amounts. What is the total vapor pressure of the

mixture and what are their mole fractions in the

vapor phase?

Prentice Hall ©2004 CHAPTER 11 SOLUTIONS AND THEIR PROPERTIES Chapter 11Slide 1.

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