ch13 properties of-solutions

Solutions

Chapter 13

Properties of Solutions

Section 13.1-13.5 are covered

Chemistry, The Central Science, 11th edition

Theodore L. Brown; H. Eugene LeMay, Jr.;

and Bruce E. Bursten

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Solutions

13.1) The Solution Process

• 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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Solutions

© 2009, Prentice-Hall, Inc.

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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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How Does a Solution Form?

As a solution forms, the solvent pulls solute

particles apart and surrounds, or solvates,

them.

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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.

The formation of a

solution can be

either exothermic or

endothermic.

Example: dissolution of NaOH in water; exothermic

Hsoln = -44.48 kJ/mol

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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 NH4NO3 in water,

heat is absorbed,

not released.

i.e. is endothermic,

Hsoln = 26.4 kJ/mol

Ammonium nitrate has been used to make

instant ice packs.

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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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Student, Beware!

Just because a substance disappears when it

comes in contact with a solvent, it doesn’t

mean the substance dissolved.

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Student, Beware!

• Dissolution is a physical change — you can get back the

original solute by evaporating the solvent.

• If you can’t, the substance didn’t dissolve, it reacted.

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3.2) Saturated Solutions and Solubility

• Saturated Solvent holds as much

solute as is possible at that

temperature.

Dissolved solute is in

dynamic equilibrium with

solid solute particles.

The amount of solute

needed to form a saturated

solution in a given quantity

of solvent at a specified temperature is known as

the Solubility of that solute.

Solute + Solvent Solution dissolve

crystallize

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Types of Solutions

• Unsaturated

Less than the

maximum amount of

solute for that

temperature is

dissolved in the

solvent.

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Types of Solutions

• Supersaturated

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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For example, the solubility of NaCl in water at 0 C is 35.7 g per 100 ml of water.

This is the maximum amount of NaCl (35.7 g) that can be dissolved in 100 ml of

water to give a stable equilibrium solution at 0 C.

If we dissolve less than 35.7 g of NaCl in 100 ml of water at 0 C, this will form

unsaturated solution.

If we dissolve greater than 35.7 g of NaCl in 100 ml of water at 0 C, this will

form supersaturated solution.

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13.3) 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.

The extent to which one substance dissolves in another depends

on the nature of both the solute and solvent (solute-solvent

interactions). It also depends on Temperature and,

at least for gases, on Pressure.

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Factors Affecting Solubility

The more similar the

intermolecular

attractions, the more

likely one substance

is to be soluble in

another.

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Glucose (which has

hydrogen bonding)

is very soluble in

water, while

cyclohexane (which

only has dispersion

forces) is not.

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• 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.

• The solubility of a

gas in a liquid is

directly proportional

to its pressure.

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Henry’s Law

Sg = kPg

where

• Sg is the solubility of

the gas;

• k is the Henry’s law

constant for that gas in

that solvent;

• Pg is the partial

pressure of the gas

above the liquid.

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Temperature

Generally, the

solubility of solid

solutes in liquid

solvents increases

with increasing

temperature.

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Temperature

• The opposite is true

of gases:

Carbonated soft

drinks are more

―bubbly‖ if stored in

the refrigerator.

Warm lakes have

less O2 dissolved in

them than cool lakes.

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13.4) Ways of

Expressing

Concentration

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Mass Percentage

Mass % of A = mass of A in solution

total mass of solution 100%

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Parts per Million and

Parts per Billion

ppm = mass of A in solution

total mass of solution 106

Parts per Million (ppm)

Parts per Billion (ppb)

ppb = mass of A in solution

total mass of solution 109

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SAMPLE EXERCISE 13.4 Calculation of Mass-Related Concentrations

(a) A solution is made by dissolving 13.5 g of glucose (C6H12O6) in 0.100 kg of water. What is the mass

percentage of solute in this solution? (b) A 2.5-g sample of groundwater was found to contain 5.4g of

Zn2+ What is the concentration of Zn2+ in parts per million?

PRACTICE EXERCISE (a) Calculate the mass percentage of NaCl in a solution containing 1.50 g of NaCl in 50.0 g of water.

(b) A commercial bleaching solution contains 3.62 mass % sodium hypochlorite, NaOCl. What is the

mass of NaOCl in a bottle containing 2500 g of bleaching solution?

Answers: (a) 2.91%, (b) 90.5 g of NaOCl

Comment: The mass percentage of water in this solution is (100 – 11.9)% = 88.1%.

(b) Analyze: In this case we are given the number of micrograms of solute. Because 1g is

1 10–6 g, 5.4g = 5.4 10–6 g.

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moles of A

total moles in solution XA =

Mole Fraction (X)

• In some applications, one needs the

mole fraction of solvent, not solute—

make sure you find the quantity you

need!

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mol of solute

Volume of solution (in Liters) M =

Molarity (M)

• You will recall this concentration

measure from Chapter 4.

• Because volume is temperature

dependent, molarity can change with

temperature.

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mol of solute

kg of solvent m =

Molality (m)

Because both moles and mass do not

change with temperature, molality

(unlike molarity) is not temperature

dependent.

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SAMPLE EXERCISE 13.5 Calculation of Molality

A solution is made by dissolving 4.35 g glucose (C6H12O6) in 25.0 mL of water at 25°C. Calculate the

molality of glucose in the solution.

PRACTICE EXERCISE What is the molality of a solution made by dissolving 36.5 g of naphthalene (C10H8) in 425 g of toluene

(C7H8)?

Answer: 0.670 m

Solution

Solve: Use the molar mass of glucose, 180.2 g/mol, to convert grams to moles:

Because water has a density of 1.00 g/mL, the mass of the solvent is

Finally, use Equation 13.9 to obtain the molality:

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Solutions

SAMPLE EXERCISE 13.6 Calculation of Mole Fraction and Molality

An aqueous solution of hydrochloric acid contains 36% HCl by mass. (a) Calculate the mole fraction of

HCl in the solution. (b) Calculate the molality of HCl in the solution.

Solve: (a) To calculate the mole fraction of HCl, we convert the masses of HCl and H2O to moles

and then use Equation 13.7:

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PRACTICE EXERCISE A commercial bleach solution contains 3.62 mass % NaOCl in water. Calculate (a) the molality and (b)

the mole fraction of NaOCl in the solution.

Answers: (a) 0.505 m, (b) 9.00 10–3

SAMPLE EXERCISE 13.6 continued

(b) To calculate the molality of HCl in the solution, we use Equation 13.9. We calculated the number of

moles of HCl in part (a), and the mass of solvent is 64 g = 0.064 kg:

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Solutions

SAMPLE EXERCISE 13.7 Calculation of Molarity

A solution contains 5.0 g of toluene (C7H8) and 225 g of benzene and has a density of 0.876 g/mL.

Calculate the molarity of the solution.

Plan: The molarity of a solution is the number of moles of solute divided by the number of liters of

solution (Equation 13.8). The number of moles of solute (C7H8) is calculated from the number of grams

of solute and its molar mass. The volume of the solution is obtained from the mass of the solution

(mass of solute + mass of solvent = 5.0 g + 225 g = 230 g) and its density.

Solve: The number of moles of solute is

The density of the solution is used to convert the mass of the solution to its volume:

Molarity is moles of solute per liter of solution:

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Solutions

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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Solutions


13.5) Colligative Properties

• Changes in colligative properties

depend only on the number of solute

particles present, not on the identity of

the solute particles.

• Among colligative properties are

– Vapor pressure lowering

– Boiling point elevation

– Melting point depression

– Osmotic pressure

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Vapor Pressure

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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Vapor Pressure

Therefore, the vapor

pressure of a solution

is lower than that of

the pure solvent.

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Raoult’s Law

PA = XAPA where

– XA is the mole fraction of solvent (A), and

– PA is the normal vapor pressure of pure

solvent (A) at that temperature.

NOTE: This is one of those times when you

want to make sure you have the vapor

pressure of the solvent.

Quantitatively, the vapor pressure of solutions containing

Nonvolatile solutes is given by Raoult’s law:

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The solution that obeys Raoult’s law is called ideal solution.

Real solutions best approximate ideal behavior when:

1- the solute concentration is low, and

2- the solute and solvent have similar molecular sizes and

Similar types of intermolecular attractions.

Example; the vapor pressure of water is 17.4 torr at 20C.

Adding glucose to water so that the solution has Xwater = 0.80 and

Xglucose = 0.20 at constant temperature, calculate the vapor pressure

Over the solution.

Pwater = Xwater Pwater = 0.80 x 17.5 = 14 torr

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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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Boiling Point Elevation

• The change in boiling

point is proportional to

the molality of the

solution:

Tb = Kb m

where Kb is the molal

boiling point elevation

constant, a property of

the solvent. Tb is added to the normal

boiling point of the solvent.

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Boiling Point Elevation

• The change in freezing

point can be found

similarly:

Tf = Kf m

• Here Kf is the molal

freezing point

depression constant of

the solvent.

Tf is subtracted from the normal

boiling point of the solvent.

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Boiling Point Elevation and

Freezing Point Depression

Note that in both

equations, T does

not depend on what the solute is, but

only on how many particles are

dissolved.

Tb = Kb m

Tf = Kf m

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Colligative Properties of

Electrolytes

Since these properties 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

However, a 1M solution of NaCl does not show

twice the change in freezing point that a 1M

solution of methanol does.

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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

Some Na+ and Cl-

reassociate for a short

time, so the true

concentration of

particles is somewhat

less than two times the

concentration of NaCl.

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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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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 are of lower solvent concentration (higher solute concentration).

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Osmotic Pressure

The pressure required to stop osmosis,

known as osmotic pressure, , is

n

V = ( )RT = 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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Osmosis in Blood Cells

• If the solute

concentration outside

the cell is greater than

that inside the cell, the

solution is hypertonic.

• Water will flow out of

the cell, and crenation

results.

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Osmosis in Cells

• If the solute

concentration outside

the cell is less than

that inside the cell, the

solution is hypotonic.

• Water will flow into the

cell, and hemolysis

results.

ch13 properties of-solutions

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