Chapter 12: Solutions and Their Properties

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Vanessa N. Prasad-PermaulCHM 1046

Valencia Community College

Introduction

1. A mixture is any intimate combination of two or more pure substances 2. Can be classified as heterogeneous or homogeneousHeterogeneous-The mixing of components is visually nonuniform and have regions of different compositionHomogenous-Mixing is uniform, same composition throughout-Can be classified according to the size of their particles as either solutions or colloids

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Solutions

Solution 1. Homogeneous mixtures2. Contain particles with diameters in the

range of 0.1–2 nm3. Transparent but may be colored4. Do not separate on standing

Colloids 1. Milk & fog 2. Diameters 2-500 nm 3. Do not separate on standing

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Types of Solutions

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Solution Formation

Solute Dissolved substance, or smaller quantity

substance

Solvent Liquid dissolved in, larger quantity

substance

Saturated solution Contains the maximum amount of solute

that will dissolve in a given solvent.

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Solution Formation

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.

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EXAMPLE 12.1: A: GIVE AN EXAMPLE OF A SOLID SOLUTION PREPARED FROM A LIQUID AND A SOLID

A dental filling made up of liquid mercury and solid silver is a solid solution

B: GIVE AN EXAMPLE OF A LIQUID SOLUTION PREPARED BY DISSOLVING A GAS IN A LIQUID

Aqueous ammonia

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EXERCISE 12.1: IDENTIFY THE SOLUTE & SOLVENT IN THE FOLLOWING SOLUTIONS.

a) 80g of chromium & 5g of molybdenum

b) 5g of MgCl2 dissolved in 1000g of H2O

c) 39% N2, 41% Ar, and the rest O2

Three Types of interactions1. Solvent-solvent2. Solvent-solute3. Solute-solute

“Like dissolves like”solutions will form when three types of interactions are similar in kind and magnitude

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Energy Changes and the Solution Process


Example NaCl and water:Ionic solid NaCl dissolve in polar solvents like water because the strong ion-dipole attractions between Na+ and Cl- ions and polar water molecules are similar in magnitude to the strong dipole-dipole attractions between water molecules and to the strong ion-ion attractions between Na+ and Cl- ions

Example Oil and waterOil does not dissolve in water because the two liquids have different kinds of intermolecular forces. Oil is not polar or an ionic solvent

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NaCl in H2O1. Ions that are less tightly held because of their position at a

corner or an edge of the crystal are exposed to water molecules

2. Water molecules will collide with the NaCl until an ion breaks free

3. More water molecules then cluster around the ion, stabilizing it by ion-dipole attractions

4. The water molecules attack the weak part of the crystal until it is dissolved

5. Ions in solution are said to be solvated they are surrounded and stabilized by an ordered shell of solvent molecules

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EXAMPLE 12.2: WOULD NAPTHALENE (C10H8) BE MORE SOLUBLE IN ETHANOL OR IN BENZENE? EXPLAIN.

Naphthalene is more soluble in benzene because nonpolar naphthalene must break the strong hydrogen bonds between ethanol molecules and replace them with weaker London forces.

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EXERCISE 12.2: WHICH OF THE FOLLOWING COMPOUNDS IS LIKELY TO BE MORE SOLUBLE IN WATER: C4H9OH or C2H9SH? EXPLAIN.


G, Free energy change1. If G is negative the process is spontaneous, and the substance is dissolved2. If G is positive the process is non-spontaneous, the substance is not dissolved3. G = H -TSH, enthalpy, heat flow in or out of the system, Hsoln heat of solution

S, entropy, disorder, Ssoln entropy of solution

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Ssoln Entropy of SolutionUsually a positive number because when you dissolve something you are increasing disorder

Hsoln Heat of Solution1. Harder to predict because it could be exothermic (- Hsoln) or endothermic (+Hsoln)2. The value of the heat of solution for a substance results from an interplay of the three kinds of interactions

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1) Solvent-solvent interactions: Energy is required (endothermic) to overcome intermolecular forces between solvent molecules because the molecules must be separated and pushed apart to make room for solute particles

2) Solute-solute interactions: Energy is required (endothermic) to overcome interactions holding solute particles together in a crystal. For an ionic solid, this is the lattice energy. Substances with higher lattice energies therefore tend to be less soluble than substances with lower lattice energies.

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3) Solvent-solute interactions: Energy is released (exothermic) when solvent molecules cluster around solute particles and solvate them. For ionic substances in water, the amount of hydration energy released is generally greater for smaller cations than for larger ones because water molecules can approach the positive nuclei of smaller ions more closely and thus bind more tightly. Hydration energy generally increases as the charge on the ion increases.

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The solute–solvent interactions are stronger than solute–solute or solvent–solvent.

Favorable process

Exothermic rxn.

Exothermic -Hsoln


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The solute–solvent interactions are weaker than solute–solute or solvent–solvent.

Unfavorable process.

Endothermic rxn

Endothermic +Hsoln


Hydration The attraction of ions for water molecules

Hydration Energy The energy associated with the attraction

between ions and water molecules

Lattice energy The energy holding ions together in a

crystal lattice

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EXAMPLE 12.3: WHICH OF THE FOLLOWINGIONS WOULD BE EXPECTED TO HAVE THE GREATER ENERGY OF HYDRATION. Mg2+ OR Al3+?

Al3+ because water molecules can approach the positive nuclei of smaller ions more closely and thus bind more tightly. Hydration energy generally increases as the charge on the ion increases.

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EXERCISE 12.3: WHICH ION HAS THE LARGER HYDRATION ENERGY, Na+ or K+ EXPLAIN.

Molarity (M)M = mole of solute / Liter of solution

Molality (m)m = moles of solute/mass of solvent (kg)

Mole Fraction (x)X = mole of component / total moles

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Units of Concentration

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EXAMPLE 12.6: GLUCOSE, C6H12O6, IS A SUGAR THAT IS IN FRUITS. IT IS ALSO FOUND IN BLOOD AND IS THE BODY’S MAIN SOURCE OF ENERGY. WHAT IS THE MOLALITY OF A SOLUTION CONTAINING 5.67g OF GLUCOSE DISSOLVED IN 25.2g OF WATER?

5.67g C6H12O6 x 1 mol C6H12O6 = 0.0315 mol C6H12O6 180.2g C6H12O6

0.0315 mol C6H12O6 = 1.25 m C6H12O6 25.2g x 1kg 1000g

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EXERCISE 12.6: TOLUENE, C6H5CH3, IS A STARTING MATERIAL FOR TNT (TRINITROTLOUENE) . FIND THE MOLALITY OF TOLUENE IN A SOLUTION THAT CONTAINS 35.6g OF TOLUENE AND 125g OF BENZENE

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EXAMPLE 12.7: WHAT ARE THE MOLE FRACTIONS OF GLUCOSE AND WATER IN A SOLUTION CONTAINING 5.67g OF GLUCOSE DISSOLVED IN 25.2g OF WATER?

5.67g C6H12O6 x 1 mol C6H12O6 = 0.0315 mol C6H12O6 180.2g C6H12O6

25.2g H2O x 1 mol H2O = 1.40 mol H2O 18.0g H2O

1.40mol + 0.0315mol = 1.432mol

MOLE FRACTION OF GLUCOSE 0.0315mol = 0.0220 x 100% = 2.20% 1.432molMOLE FRACTION OF WATER 1.40mol = 0.978 x 100% = 97.8% 1.432mol

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EXERCISE 12.7: CALCULATE THE MOLE FRACTIONS OF TOLUENE AND BENZENE IN THE SOLUTION FROM EXERCISE 12.6

Units of Concentration

Mass Percent (mass %)Mass % = (mass of component / total mass of sol’n) 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

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EXAMPLE 12.5: HOW WOULD YOU PREPARE 425g OF AN AQUEOUS SOLUTION CONTAINING 2.40% BY MASS OF SODIUM ACETATE, NaC2H3O2?

MASS OF SOLUTE: 0.0240 x 425g = 10.2gMASS OF WATER: 425g – 10.2g = 414.8 = 415gMASS OF SOLUTION: 425g

YOU WOULD PREPARE THIS SOLUTION BY DISSOLVING 10.2g OF NaC2H3O2 IN 415g OF WATER FOR A TOTAL MASS OF 425g.

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EXERCISE 12.5: AN EXPERIMENT CALLS FOR 35.0g OF HYDROCHLORIC ACID THAT IS 20.2% HCl BY MASS. HOW MANY GRAMS OF HCl IS THIS? HOW MANY GRAMS OF WATER IS THIS?

Some Factors Affecting Solubility Solubility

The amount of solute per unit of solvent needed to form a saturated solution

MiscibleMutually soluble in all proportions

Effect of Temperature on Solubility

1. Most solid substances become more soluble as temperature rises

2. Most gases become less soluble as temperature

rises

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Some Factors Affecting Solubility

Effect of Pressure on Solubility1. No effect on liquids or solids2. The solubility of a gas in a liquid at a given temperature is directly proportional to the partial pressure of the gas over the solution, @ 25°C

Henry’s Law solubility = k x P

k = constant characteristic of specific gas, mol/Latm

P = partial pressure of the gas over the sol’n 33

Some Factors Affecting Solubility

a) Equal numbers of gas molecules escaping liquid and returning to liquid b) Increase pressure, increase # of gas molecules

returning to liquid, solubility increasesc) A new equilibrium is reached, where the #’s of

escaping = # of returning

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EXAMPLE 12.4: 27g OF ACETYLENE (C2H2) DISSOLVES IN 1L OF ACETONE AT 1.0atm. IF THE PARTIAL PRESSURE OF ACETYLENE IS INCREASED TO 12atm, WHAT IS THE SOLUBILITY OF ACETONE?

S2 = P2S1 P1

S2 = 12atm 27g C2H2/L acetone 1.0atm

S2 = 27g C2H2 x 12atm = 3.2x102g C2H2/L acetone L acetone 1.0atm

320g of acetylene will dissolve in 1L of acetone @ 12atm

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EXERCISE 12.4: A LITER OF WATER @ 250C DISSOLVES 0.0404g OF O2 WHEN THE PARTIAL PRESSURE OF THE OXYGEN IS 1.00atm. WHAT IS THE SOLUBILITY OF OXYGEN FROM AIR, IN WHICH THE PARTIAL PRESSURE OF O2 IS 159mmHg?

Physical Behavior of Solutions: Colligative

Properties Colligative properties

Properties that depend on the amount of a dissolved solute but not its chemical identity

There are four main colligative properties:1. Vapor pressure lowering2. Freezing point depression3. Boiling point elevation4. Osmotic pressure

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Physical Behavior of Solutions: Colligative Properties

In comparing the properties of a pure solvent with those of a solution…

1. Vapor pressure of sol’n is lower

2. Boiling point of sol’n is higher

3. Freezing point of sol’n is lower

4. Osmosis, the migration of solvent molecules through a semipermeable membrane, occurs when solvent and solution are separated by the membrane

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Vapor-pressure Lowering of Solutions: Raoult’s

Law 1. A liquid in a closed container is in equilibrium with its

vapor and that the amount of pressure exerted by the vapor is called the vapor pressure.

2. When you compare the vapor pressure of a pure solvent with that of a solution at the same temperature the two values are different.

3. If the solute is nonvolatile and has no appreciable vapor pressure of its own (solid dissolved) the vapor pressure of the solution is always lower that that of the pure solvent.

4. If the solute is volatile and has a significant vapor pressure (2 liquids) the vapor pressure of the mixture is

intermediate between the vapor pressures of the two pure liquids.

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Solutions with a Nonvolatile Solute

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.

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Solutions with a Nonvolatile Solute!!!

Raoult’s Law Psoln = Psolv · Xsolv

Psoln = vapor pressure of the solution

Psolv = vapor pressure of the pure solvent

Xsolv = mole fraction of the solvent in the solution

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Raoult’s Law applies to only Ideal solutions

1. Law works best when solute concentrations are low and when solute and solvent particles have similar intermolecular forces.

2. If intermolecular forces between solute particles and solvent molecules are weaker than solvent molecules alone, solvent molecules are less tightly held, vapor pressure is higher than Raoult predicts

3. If intermolecular forces between solute and solvent are stronger than solvent alone, solvent molecules are more tightly held and the vapor pressure is lower than predicted

4. No Van’t Hoff factor!!! is a measure of the effect of a solute upon colligative properties

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http://en.wikipedia.org/wiki/Colligative_properties

Solutions with a Nonvolatile Solute

Close-up view of part of the vapor pressure curve

for a pure solvent and a solution of a nonvolatile

solute. Which curve represents the pure solvent,

and which the solution?

Why?

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The lower vapor pressure of a sol’n relative to that of a pure solvent is due to the difference in their entropies of vaporization, Svap. Because the entropy of the solvent in a sol’n is higher to begin with, Svap is smaller for the sol’n than for the pure solvent. As a result vaporization of the solvent from the sol’n is less favored (less negative Gvap), and the vapor pressure of the solution is lower.

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Solutions with a Volatile Solute!! Ptotal = PA + PB

Ptotal = (P°A · XA) + (P°B · XB)

P°A = vapor pressure of pure AXA = mole fraction of AP°B = vapor pressure of pure BXB = mole fraction of B

Ptotal should be intermediate to A & B 45

Close-up view of part of the vapor pressure curves fortwo pure liquids and a mixture of the two. Whichcurves represent the mixture?

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1) Red

2) Green

3) Blue

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EXAMPLE 12.12: CALCULATE THE VAPOR-PRESSURE WHEN 5.67g OF GLUCOSE IS DISSOLVED 25.2g OF WATER @ 25oC. THE VAPOR PRESSURE OF WATER @ 25oC IS 23.8mmHg. WHAT IS THE VAPOR PRESSURE OF THE SOLUTION?

P = PAOXB = 23.8 mmHg x 0.0220 =

0.524 mmHg

THE VAPOR PRESSURE OF THE SOLUTION IS:PA

= PAO - P

= (23.8mmHg – 0.524mmHg) = 23.3mmHg

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EXERCISE 12.12: NAPTHALENE IS USED TO MAKE MOTHBALLS. A SOLUTION IS MADE BY DISSOLVING 0.515g OF NAPTHALENE IN 60.8g OF CHLOROFORM. CALCULATE THE VAPOR PRESSURE LOWERING OF CHLOROFORM @ 20oC FROM NAPTHALENE. THE VAPOR PRESSURE OF CHLOROFORM @ 20oC IS 156mmHg. NAPTHALENE CAN BE ASSUMED TO BE NON-VOLATILE COMPARED WITH CHLOROFORM. WHAT IS THE VAPOR PRESSURE OF THE SOLUTION?

Boiling Point Elevation and Freezing Point Depression of

Solutions

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Boiling Point Elevation and Freezing Point Depression of Solutions

1. Red line is pure solvent2. Green line solution of nonvolatile solute3. Vapor pressure of sol’n is lower4. Temp at which vapor pressure = 1 atm for sol’n is

higher5. Boiling point of sol’n is higher by Tb

6. Liquid/vapor phase transition line is lower for sol’n7. Triple point temp is lower for sol’n8. Solid/liquid phase transition has shifted to a lower

temp.9. The freezing point of the sol’n is lower by Tf 50


Tb = Kb · m

Tf = Kf · m

Kb = molal boiling-point elevation constantKf = molal freezing-point depression constantm = molalityNO Van’t Hoff Factor!!

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The higher boiling point of a solution relative to that of a pure

solvent is due to a difference in their entropies of vaporization,

Svap. Because the solvent in a solution has a higher entropy to

begin with, Svap is smaller for the solution than for the pure

solvent. As a result, the boiling point of the solution Tb is higher

than that of the pure solvent.

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The lower freezing point of a solution relative to that of a

pure solvent is due to a difference in their entropies of fusion,

Sfusion. Because the solvent in a solution has a higher

entropy level to begin with, Sfusion is larger for the solution

than for the pure solvent. As a result the freezing point of

the solution Tf is lower than that of the pure solvent.

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EXAMPLE 12.13: AN AQUEOUS SOLUTION IS 0.0222 m GLUCOSE. WHAT ARE THE BOILING POINT AND THE FREEZING POINT OF THIS SOLUTION?

Tb = Kbcm = 0.512oC/m x 0.0222 m = 0.0114oCTf = Kfcm = 1.86oC/m x 0.0222 m = 0.0413oC

THE BOILING POINT OF THE SOLUTION IS 100.000oC + 0.0114oC = 100.011oC

THE FREEZING POINT OF THE SOLUTION IS 0.000oC – 0.0413oC = -0.041oC

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EXERCISE 12.13: HOW MANY GRAMS OF ETHYLENE GLYCOL MUST BE ADDED TO 37.8g OF WATER TO GIVE A FREEZING POINT OF -0.150oC?

Osmosis and Osmotic Pressure 1. Semipermeable membranes allow water or

other small molecules to pass through, but they block the passage of large solute molecules or ions.

2. When a solution and a pure solvent are separated by the right kind of semipermeable membrane, solvent molecules pass through the membrane in a process known as osmosis.

3. Passage of solvent through the membrane takes place in both directions

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Osmosis and Osmotic Pressure

4. Passage from the pure solvent side to the solution side is more favored and faster.

5. The amount of liquid on the pure solvent side decreases

6. The amount of liquid on the solution side increases

7. The concentration of the solution decreases

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Osmotic Pressure

1. The amount of pressure necessary to achieve equilibrium

2. = MRT = osmotic pressureM = molarityR = gas constant, .08206 L atm/K molT = temperature in kelvins

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Some uses of Colligative Properties

1. The most important use of colligative properties in the laboratory is for determining the molecular mass of an unknown substance.

2. Any of the four colligative properties can be used but using osmotic pressure gives the most accurate results

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If you have to calculate the molar mass of a compound

which of the following will give you the most accurate

results?

1) Osmotic pressure2) Vapor pressure lowering3) Boiling point elevation4) Freezing point depression

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EXAMPLE 12.8: AN AQUEOUS SOLUTION IS 0.120 m GLUCOSE. WHAT ARE THE MOLE FRACTIONS OF EACH COMPONENT IN THE SOLUTION?

0.120 m = 0.120mol IN 1.00kg OF WATER1.00kg H2O x 1000g x 1 mol = 55.6mol H2O 1kg 18.0 g H2O

MOLE FRACTION OF GLUCOSE: 0.120 mol = 0.00215

(0.120mol + 55.6mol)

MOLE FRACTION OF WATER: 55.6mol = 0.998 (0.120mol + 55.6mol)

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EXERCISE 12.8: A SOLUTION IS 0.120 m METHANOL DISSOLVED IN ETHANOL. CALCULATE THE MOLE FRACTIONS OF CH3OH AND CH3CH2OH IN THE SOLUTION.

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EXAMPLE 12.9: A SOLUTION IS 0.150 mol FRACTION GLUCOSE, AND 0.850 mol FRACTION WATER. WHAT IS THE MOLALITY OF GLUCOSE IN THE SOLUTION?

0.850 mol x 18.0 g H2O = 15.3 g H2O (0.0153kg H2O) 1 mol H2O

0.150 mol C6H12O6 = 9.80 m C6H12O60.00153kg SOLVENT

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EXERCISE 12.9: A SOLUTION IS 0.250 mol FRACTION METHANOL AND 0.750 mol FRACTION ETHANOL. WHAT IS THE MOLALITY OF METHANOL IN THE SOLUTION?

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EXAMPLE 12.10: AN AQUEOUS SOLUTION IS 0.273 m KCl. WHAT IS THE MOLAR CONCENTRATION OF KCl? THE DENSITY OF THE SOLUTION IS 1.011 x 103g/L.

0.273 mol KCl x 74.6g KCl = 20.4g KCl 1 mol KCl

1.000 x 103g + 20.4g = 1.020 x 103g

1.020 x 103g x 1L = 1.009L 1.011x103g

0.273mol KCl = 0.271 M KCl 1.009L solution

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EXERCISE 12.10: UREA (NH2)2CO IS USED AS FERTILIZER. WHAT IS THE MOLAR CONCENTRATION OF AN AQUEOUS SOLUTION THAT IS 3.42 m UREA? THE DENSITY OF THE SOLUTION IS 1.045 g/mL.

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EXAMPLE 12.11: AN AQUEOUS SOLUTION IS 0.907 M Pb(NO3)2. WHAT IS THE MOLALITY? THE DENSITY OF THE SOLUTION IS 1.252g/mL.

1.000 L x 1000mL x 1.252 g = 1.252 x 103g 1L 1mL

0.907 mol Pb(NO3)2 x 331.2 g Pb(NO3)2 = 3.00 x 102 Pb(NO3)2 1 mol

1.252 x 103 g – 3.00 x 12 g = 9.52 x 102g (0.952kg)

0.907 mol Pb(NO3)2 = 0.953 m Pb(NO3)2 0.952kg solvent

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EXERCISE 12.11: AN AQUEOUS SOLUTION IS 2.00 M UREA. THE DENSITY IS 1.029 g/mL. WHAT IS THE MOLAL CONCENTRATION OF UREA IN THE SOLUTION?

Example 1

Arrange the following in order of their expected increasing solubility in water:

Br2, KBr, C7H8

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Example 2

1. Na+

2. Cs+

3. Li+

4. Rb+

1. Mg2+

2. Na+

3. Li+

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Which would you expect to have the larger (more negative) hydration energy?

Example 3

What is the mass % concentration of a saline sol’n prepared by dissolving 1.00 mol of NaCl in 1.00 L of water? DensityH2O=1.00 g/mL MMNaCl = 58.443 g/mol

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Example 4

Assuming that seawater is an aqueous solution of NaCl what is its molarity? The density of seawater is 1.025 g/mL at 20C and the NaCl concentration is 3.50 mass %

Assume 1 L to make easier, 1000 mL

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Example 5

What is the molality of a solution prepared by dissolving 0.385 g of cholesterol, C27H46O in 40.0 g of chloroform, CHCl3? What is the mole fraction of cholesterol in the solution?

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Example 6

What mass in grams of a 0.500 m solution of sodium acetate, CH3CO2Na, in water would you use to obtain 0.150 mol of sodium acetate?

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Example 7

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/mol. What is the molarity of the solution?

Assume 1 kg

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Example 8

The density at 20°C of a 0.500 M solution of acetic acid in water is 1.0042 g/mL. What is the concentration of this solution in molality? The molar mass of acetic acid, CH3CO2H, is 60.05 g/mol.

Assume 1 L

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Example 9

Which of the following will become less soluble in water as the temperature is increased?

1) NaOH(s)

2) CO2(g)

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Example 10

The solubility of CO2 in water is 3.2 x 10-2 M @ 25°C

and 1 atm pressure. What is the Henry’s-Law constant

for CO2 in mol/L atm?

solubility = k x P

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Example 11

What is the vapor pressure (in mm Hg) of a solution

prepared by dissolving 5.00 g of benzoic acid

(C7H6O2) in 100.00 g of ethyl alcohol (C2H6O) at

35°C? The vapor pressure of the pure ethyl alcohol

at 35°C is 100.5 mm Hg

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Example 12

What is the vapor pressure ( in mm Hg) of a sol’n

prepared by dissolving 25.0 g of ethyl alcohol

(C2H5OH) in 100.0 g of water at 25°C? The vapor

pressure of pure water is 23.8 mm Hg and the vapor

pressure of ethyl alcohol is 61.2 mm Hg at 25°C

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Example 13

a) Is the boiling point of the second pure liquid higher or lower than that of the first liquid?

b) On the diagram where is the approximate position of the second pure liquid?

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The following phase diagram shows part of the vapor pressure curves for a pure liquid (green curve) and a solution (red curve) of the first liquid with a second volatile liquid (not shown)

Example 14

What is the normal boiling point in °C of a solution

prepared by dissolving 1.50 g of aspirin (C9H8O4) in

75.00 g of chloroform (CHCl3)? The normal boiling

point of chloroform is 61.7 °C and Kb of chloroform

is 3.63 °C kg/mol

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Example 15

What osmotic pressure in atm would you expect for a solution of 0.125 M C6H12O6 that is separated from pure water by a semipermeable membrane at 310 K?

= MRT

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Example 16

A solution of unknown substance in water at 300 K gives rise to an osmotic pressure of 3.85 atm. What is the molarity of the solution?

= MRTM = /RT

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Example 17

What is the molar mass of sucrose if a solution prepared by dissolving 0.822 g of sucrose in water and diluting to a volume of 300.0 mL has an osmotic pressure of 149 mm Hg at 298 K?

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Chapter 12: Solutions and Their Properties