Solutions

•

Solutions

are homogeneous mixtures of two or more pure substances.

• In a solution, the

solute

is dispersed uniformly throughout the

solvent

.

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The intermolecular forces between solute and solvent particles must be strong enough to compete with those between solute particles and those between solvent particles.

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Click Here to play movie:How Does a Solution Form?

As a solution forms, the solvent pulls solute particles apart and surrounds, or

solvates

, them.

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Solutions

How Does a Solution Form?

If an ionic salt is soluble in water, it is because the ion – dipole interactions are strong enough to overcome the lattice energy of the salt crystal.

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Energy Changes in Solution

• Simply put, three processes affect the energetics of solution: – Separation of solute particles, – Separation of solvent particles, – New interactions between solute and solvent.

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Energy Changes in Solution

The enthalpy change of the overall process depends on 

H

for each of these steps.

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Why Do Endothermic Processes Occur?

Things do not tend to occur spontaneously (i.e., without outside intervention) unless the energy of the system is lowered.

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Why Do Endothermic Processes Occur?

Yet we know that in some processes, like the dissolution of NH 4 NO 3 in water, heat is absorbed, not released.

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Enthalpy Is Only Part of the Picture

The reason is that increasing the disorder or randomness (known as

entropy

) of a system tends to lower the energy of the system.

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Enthalpy Is Only Part of the Picture

So even though enthalpy may increase, the overall energy of the system can still decrease if the system becomes more disordered.

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Types of Solutions

• Saturated – In a

saturated

solution, the solvent holds as much solute as is possible at that temperature.

– Dissolved solute is in dynamic equilibrium with solid solute particles.

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Types of Solutions

• Unsaturated – If a solution is

unsaturated

, less solute than can dissolve in the solvent at that temperature is dissolved in the solvent.

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Solutions

Types of Solutions

• Supersaturated – In

supersaturated

solutions, the solvent holds more solute than is normally possible at that temperature.

– These solutions are unstable; crystallization can usually be stimulated by adding a “seed crystal” or scratching the side of the flask.

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Factors Affecting Solubility

• Chemists use the axiom “like dissolves like.” – Polar substances tend to dissolve in polar solvents.

– Nonpolar substances tend to dissolve in nonpolar solvents.

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Exercise 1

• • Decide whether liquid hexane (C6H14) or liquid methanol (CH3OH) is the more appropriate solvent for the substances grease (C20H42) and potassium iodide (KI).

Solutions

Factors Affecting Solubility

The more similar the intermolecular attractions, the more likely one substance is to be soluble in another.

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Factors Affecting Solubility

Glucose (which has hydrogen bonding) is very soluble in water, while cyclohexane (which only has dispersion forces) is not.

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Factors Affecting Solubility

• Vitamin A is soluble in nonpolar compounds (like fats).

• Vitamin C is soluble in water.

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Gases in Solution

• In general, the solubility

of gases in water increases with increasing mass

.

• Larger molecules have stronger dispersion forces.

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Gases in Solution

• The solubility of liquids and solids does not change appreciably with pressure.

• Click here to play movie:But the solubility of a gas in a liquid is directly proportional to its pressure.

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Henry’s Law

S g

=

kP g

• • • where

S g

is the solubility of the gas,

k

is the Henry’s Law constant for that gas in that solvent, and

P g

is the partial pressure of the gas above the liquid.

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Exercise 2

A certain soft drink is bottled so that a bottle at 25 ° C contains CO2 gas at a pressure of 5.0 atm over the liquid. Assuming that the partial pressure of CO2 in the atmosphere is 4.0 × 10−4 atm, calculate the equilibrium concentrations of CO2 in the soda both before and after the bottle is opened. The Henry’s law constant for CO2 in aqueous solution is 0.031 mol/L • atm at 25 ° C.

before = 0.16 mol/L after = 1.2

×

10 —5 mol/L

Solutions

Temperature

• Generally, the solubility of solid solutes in liquid solvents increases with increasing temperature. Most solution formation is endothermic — LeChatelier’s Principle ). Solubility generally decreases with

Temperature

• The opposite is true of gases.

– Carbonated soft drinks are more “bubbly” if stored in the refrigerator.

– Warm lakes have less O 2 dissolved in them than cool lakes.

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Temperature/Solubility of gases

• The opposite is true of gases.

–

Carbonated

soft drinks are more “bubbly” if stored in the refrigerator.

The solubility of a gas in water

always decreases with increasing temperature

.

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Ways of Expressing Concentrations of Solutions

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Mass Percentage

Mass % of

A

= mass of

A

in solution total mass of solution  100 Solutions © 2012 Pearson Education, Inc.

Mole Fraction (

X

)

X A

= moles of

A

total moles of all components • In some applications, one needs the mole fraction of

solvent

, not solute — make sure you find the quantity you need!

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Molality (

m

)

m

= moles of solute kilograms of solvent Since both moles and mass do not change with temperature, molality (unlike molarity) is

not

temperature dependent.

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Changing Molarity to Molality

If we know the density of the solution, we can calculate the molality from the molarity, and vice versa.

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Colligative Properties

• Changes in

colligative properties

depend only on the

number identity

of solute particles present, not on the of the solute particles.

• Among colligative properties are – Vapor-pressure lowering – Boiling-point elevation – Melting-point depression – Osmotic pressure – Click Here for the Chemguy’s discussion of colligative properties

–Stop here

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Solutions

How many moles of ions in 250.0 ml of .500 M Al

2

(SO

4

)

3 Solutions

1. Vapor Pressure Lowering

Because of solute –solvent intermolecular attraction, higher concentrations of

nonvolatile solutes

make it harder for solvent to escape to the vapor phase.

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1. Vapor Pressure Lowering

OR A more complicated (more favored), entropy based explanation. When a solute is added to a pure solvent, the entropy (disorder) dramatically increases. This is the preferred situation (since positive  S helps  G be negative). Prior to the solute being added, the solvent could only increase the disorder by evaporating (becoming vapor). With the solute added the entropy has already been dramatically increased, and there is less incentive for the vaporization process.

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Vapor Pressure

Therefore, the vapor pressure of © 2012 Pearson Education, Inc.

Solutions

The vapor pressure of a solution is

directly proportional to the mole fraction of solvent

present. If the solute ionizes , the number of ions further affects (lowers) the vapor pressure.

P solution = observed vapor pressure of the solvent in the solution

χ solvent = mole fraction of solvent

Po

solvent vapor pressure of pure solvent

i = van’t Hoff factor (moles of electrolyte must be multiplied by this)

number of moles particles in solution/number of moles particles dissolved Solutions

Changes in vapor pressure are governed by

Raoult’s Law.

Fall into 2 categories:

1. Those where the solute is non-volatile, 2.and those where the solute is volatile, i.e.

those where the solution has two volatile components Solutions

Exercise 5 Calculating the Vapor Pressure of a Solution

Calculate the expected vapor pressure at 25 ° C for a solution prepared by dissolving 158.0 g of common table sugar (sucrose, molar mass = 342.3 g/mol) in 643.5 cm3 of water. At 25 ° C, the density of water is 0.9971 g/cm3 and the vapor pressure is 23.76 torr.

= 23.46 torr

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Boiling-Point Elevation

• The boiling point of a solution is reached when the external atmospheric pressure is equal to its vapor pressure. Since addition of a solute lowers the vapor pressure of a solution, it must, by definition, raise the boiling point.

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Boiling Point Elevation

• For

non-electrolyte solutions (non ionic), the change is summarized in the expression

• T b = K b m   Tb = change in boiling point of the solvent, K b = molal boiling point constant for the solvent and • m = molality of the solute Solutions

Freezing point depression

• Freezing points are lowered by the addition of non-volatile solutes.

• Again, this can be explained in one of two ways; • (i) As a physical phenomenon (less favored), where solute particles actually get in the way of the solid structure being formed. If the solidification (freezing) process is hindered, more heat must be removed from the solution, and hence the freezing • point is lowered.

Freezing point depression

Freezing point depression

• OR • (ii) As an entropy based phenomenon (more favored), where the solution has more disorder than the pure solvent so needs more energy removed from it in order to create the more ordered solid.



T f

=

K f



m

Boiling-Point Elevation and Freezing-Point Depression

Nonvolatile solute – solvent interactions also cause solutions to have higher boiling points and lower freezing points than the pure solvent.

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• •

EXAMPLE: Calculate the freezing point and boiling point of a solution of 100. g ethylene glycol (C2H6O2) in 900. g of water.

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Osmosis

• Some substances form

semipermeable membranes

, allowing some smaller particles to pass through, but blocking other larger particles.

• In biological systems, most semipermeable membranes allow water to pass through, but solutes are not free to do so.

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Osmosis

In osmosis, there is net movement of solvent from the area of

higher solvent concentration

(

lower solute concentration

) to the area of

lower solvent concentration

(

higher solute concentration

).

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Osmotic Pressure

The pressure required to stop osmosis, known as

osmotic pressure

,  , is  =

( )

RT V

=

MRT

where

M

is the molarity of the solution.

If the osmotic pressure is the same on both sides of a membrane (i.e., the concentrations are the same), the solutions are

isotonic

.

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Solutions

•

Exercise 9 Determining Molar Mass from Osmotic Pressure

To determine the molar mass of a certain protein, 1.00 × 10−3 g of it was dissolved in enough water to make 1.00 mL of solution. The osmotic pressure of this solution was found to be 1.12 torr at 25.0

° C. Calculate the molar mass of the protein.

= 1.66

×

10 4 g/mol

Solutions

Electrolyte solutions and the van’t Hoff factor

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Colligative Properties of Electrolytes

Since the colligative properties of electrolytes depend on the number of particles dissolved, solutions of electrolytes (which dissociate in solution) should show greater changes than those of nonelectrolytes.

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Colligative Properties of Electrolytes

For example, one unit of MgCl2 splits into three particles when it is dissolved in solution (one Mg2+ ion and two Cl- ions). Since we are dealing with colligative properties (those that depend on the number of particles present), this should produce a freezing point depression three times greater than a non-electrolyte solute that does not produce a greater number of particles when dissolved, for example, molecular sucrose, C6H12O6.

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Colligative Properties of Electrolytes

• In order to account for this, the van’t Hoff factor (i) is added to the expressions already discussed.

• • • T b 

T f

= i K b = i

K f

 m

m

 = ( n/v )

RT

= i

MR T

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van’t Hoff Factor

One mole of NaCl in water does not really give rise to two moles of ions.

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van’t Hoff Factor

• Reassociation is more likely at higher concentration.

• Therefore, the number of particles present is concentration dependent.

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Colloids

Suspensions of particles larger than individual ions or molecules, but too small to be settled out by gravity, are called

colloids

.

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Tyndall Effect

• Colloidal suspensions can scatter rays of light.

• This phenomenon is known as the

Tyndall effect

.

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Colloids in Biological Systems

Some molecules have a polar,

hydrophilic

(

water-loving

) end and a nonpolar,

hydrophobic

hating

) end.

(

water-

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Colloids in Biological Systems

These molecules can aid in the emulsification of fats and oils in aqueous solutions.

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