David A. Katz Department of Chemistry Pima Community College

(1)

Solutions

David A. Katz

Department of Chemistry Pima Community College

(2)

A A solution is asolution is a O tit t i ll O tit t i ll A A solution is a solution is a

HOMOGENEOUS mixture mixture of 2 or more substances in of 2 or more substances in a single phase

a single phase

One constituent is usually One constituent is usually regarded as the

regarded as the SOLVENT

(usually water) and the others and the others a single phase.

a single phase.

as

(3)

Solutions

• In a solution, the

solute

is dispersed

uniformly throughout the

solvent

(4)

Solutions:

Gases mixed

with gases

(5)

S l ti

Solutions:

Gas mixed

(6)

Solutions:

Liquid mixed

with liquid

(7)

Solutions:

Solid mixed

with liquid

(8)

Solutions:

Gas mixed

with solid

Photomicrographs Photomicrographs of Hydrogen on palladium

(9)

Solutions:

Liquid mixed

with solid

(a mercury amalgam with gold)

(10)

Solutions:

Solid mixed with solid

(alloys)

(11)

Intermolecular Forces

Wh d

b t

di

l ?

Why does a substance dissolve?

The intermolecular

forces between solute

and solvent particles

must be strong enough

to compete with those

between solute

particles and those

between solvent

(12)

Intermolecular

F

stronger

Forces:

Why does a

substance dissolve?

The intermolecular The intermolecular

forces between solute and solvent particles

t b t

must be strong

enough to compete with those between solute particles and those between solvent particles

particles.

(13)

stronger

Intermolecular

Forces:

Why does a

substance dissolve?

(continued) weaker

(14)

How Does a Solution Form?

In this example, we have an ionic solid, NaCl, and a polar solvent, H₂O.

a polar solvent, H₂O.

The solution forms because the solvent pulls

solute particles apart and surrounds, or solvates, them. In water this is called hydration.

Sol te (NaCl) in The sol te is H drated ions in

Solute (NaCl) in water The solute is dissolving Hydrated ions in solution

(15)

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 strong enough to overcome the lattice energy of the salt crystal.

The ions are hydrated – that prevents the ions from

prevents the ions from

reforming the crystal lattice under normal conditions.

(16)

Solutions

• Just because a substance disappears when it comes

in contact with a solvent, it doesn’t mean the substance dissolved

substance dissolved.

• Dissolution is a physical change — you can get back the original solute by evaporating the solvent.

(17)

Water as a Solvent

• How water dissolves molecular

How water dissolves molecular

compounds:

When the nonpolar part of an organic

–

When the nonpolar part of an organic

molecule is considerably larger than the

polar part, the molecule no longer

p

,

g

dissolves in water.

For example ethanol, CH₃CH₂OH is soluble in water but butanol CH₃CH₂CH₂CH₂OH is not

(18)

Factors Affecting Solubility

g

y

• Chemists use the axiom “like dissolves like”:

– Polar and ionic substances tend to dissolve in polar

solvents.

(19)

Factors Affecting Solubility

g

y

Solubility in water

decreases as the

nonpolar end of

Methanol

p

the alcohol

molecules

increases

Butanol Heptanol

(20)

Factors Affecting Solubility

The more similar the intermolecular

attractions the more likely one substance is

attractions, the more likely one substance is

to be soluble in another.

(21)

Factors Affecting Solubility

The more similar the intermolecular

attractions the more likely one substance is

attractions, the more likely one substance is

to be soluble in another.

(22)

Factors Affecting Solubility

Glucose (which

has hydrogen

bonding) is very

soluble in

water, while

Cyclohexane

y

(which only has

dispersion

(23)

Factors Affecting Solubility

• Vitamin A is soluble in nonpolar compounds

(like fats).

(24)

Intermolecular Forces

(25)

Energy Changes in Solution

• Three processes affect

the energetics of the

solution process:

1 Separation of solute 1. Separation of solute particles 2. Separation of solvent particles 3. Interactions (attraction) between solute and

between solute and solvent

Note: Steps 1 and 2 are sometimes combined as a single step

(26)

Energy Changes in Solution

The enthalpy

change of the

g

solution

process

d

depends on

H for each

of these

steps.

Exothermic Endothermic

(27)

The Exothermic Solution

P

Process

We can break down the process into separate steps:

1 Ions break free from 1. Ions break free from

crystal lattice, ∆H₁

2. Intermolecular forces between solvent, ∆H₂ molecules are “broken”

3 H d ti ( l ti ) f

3. Hydration (solvation) of the ions by water

molecules, ∆H₃₃

(28)

The Exothermic Solution

P

Process

Energy (enthalpy) of hydration (solvation) of the ions Energy (enthalpy) of hydration (solvation) of the ions by water molecules, increases with increasing

(29)

The Exothermic

Solution

Process

(30)

Calculating the Enthalpy of an

E

th

i S l ti

Exothermic Solution

In this example, KF is dissolved in water. Step 2, the

ti f l t l l i i t d i t th

separation of solvent molecules, is incorporated into the lattice energy

(31)

The Endothermic Solution

P

Process

The process occurs

in the same steps

as an exothermic

process

The energies are

different:

∆H

₃

 ∆H

₁

+ ∆H

₂

(32)

Enthalpy Is Only Part of the Picture

The reason is that

increasing the disorder or

randomness (known as

) f

entropy

) of a system tends

to lower the energy of the

system

system.

So even though enthalpy

may increase the overall

may increase, the overall

energy of the system can

still decrease if the system

y

becomes more disordered.

(33)

Types of Solutions

y

Saturated

Solvent holds as much solute as is possible at that temperature.

To insure saturation, a small amount of solute remains undissolved on of solute remains undissolved on the bottom of the container

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

(34)

Types of Solutions

yp

• Unsaturated

L th th i

Less than the maximum amount of solute for that temperature is dissolved in the solvent.

The amount of solute in the solution can vary from a

solution can vary from a small amount to almost saturated

(35)

Types of Solutions

• Supersaturated

Solvent holds more solute than is normally

Solvent holds more solute than is normally possible at that temperature.

These solutions are unstable; crystallization

ll b i l d b ddi “ d

can usually be stimulated by adding a “seed crystal” or scratching the side of the flask.

(36)

Gases in Solution

• In general, the

l bilit

f

solubility of gases

in water increases

with increasing

mass.

• Larger molecules

Larger molecules

have stronger

(37)

Gases in Solution

The sol bilit of liq ids

• The solubility of liquids

and solids does not

change appreciably

with pressure.

• The solubility of a gas

The solubility of a gas

in a liquid is directly

proportional to its

pressure in contact

with the liquid.

(38)

Henry’s Law

Henry s Law

S

_g

= kP

_g

where

• S

_g

is the solubility of

the gas;

• k

is the Henry’s law

t

t f

th t

constant for that gas

in that solvent;

P

i th

ti l

• P

_g

is the partial

pressure of the gas

above the liquid

(39)

Temperature

Generally, the

l bilit

f

lid

solubility of solid

solutes in liquid

solvents

increases with

increasing

g

temperature.

(40)

Temperature

• The opposite is

true of gases:

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

W l k h

Warm lakes have less O₂ dissolved in them than cool lakes.

(41)

Some Henry’s Law Constants for Common Gases at 25°C G k ( l/ ) Gas k_H (mol/atm) H₂ 7.8 x 10‐4 N₂ 6.1 x 10‐4 O 1 3 x 10‐3 O₂ 1.3 x 103 Cl₂ 9.3 x 102 Br₂ 7.1 x 10‐1 I₂₂ 3.1 He 3.7 x 10‐4 Ne 4.5 x 10‐4 Ar 1.4 x 10‐3 NH₃ 6.1 x 101 CO₂ 3.4 x 10‐2 NO₂ 4.0 x 10‐2 SO 1 2 SO₂ 1.2 HCl 2.0 x 101 CH₄ 1.4 x 10‐3 C₂H₆ 1.9 x 10‐3

Reference: Sander, Rolf, Compilation of Henry’s Law Constants for Inorganic and Organic Species of

Potential Importance in Environmental Chemistry, 1999

(42)

Gases in Solution

The Fizz-Keeper

Claims to keep “soft

d i k f i

drinks from going flat!”

Does it work? Does it work?

Why?

(43)

How does

the solubility

y

of a gas

change with

g

(44)

Expressing

p

g

Concentrations of

Solutions

(45)

Percent, %

Can be expressed as:

Percent by mass, %_(m/m)

Percent by volume, %_(v/v) for solutions of liquids Percent mass-volume, %_(m/v) for solids in liquids

% of A =

amount of A in solution

t t l

t f

l ti

 100

total amount of solution

(46)

Parts per Million and

P t

Billi

Parts per Billion

mass of A in solution

Parts per Million (ppm)

ppm =

mass of A in solution

total mass of solution

 10

6

Parts per Billion (ppb)

ppb =

mass of A in solution

(47)

Molarity (M)

mol of solute

Molarity (M)

mol of solute

L of solution

M =

An alternate equation is An alternate equation is

1000

mL

g







1000

mL L solute solute solution

g

M

MW

mL

• Note: Because volume is temperature dependent,

M l it h ith t t

(48)

Mole Fraction (X)

moles of A

Mole Fraction (X)

moles of A

total moles in solution

X

_A

=

For a solution of two or more components, A,

B etc

B, etc…

(49)

Molality (m)

mol of solute

Molality (m)

mol of solute

kg of solvent

m =

An alternate equation is An alternate equation is

1000

g kg solute

g





kg solute solute solvent

g

m

MW

g

Note: Because both moles and mass do not change with temperature, molality (unlike molarity) is not

t t d d t

(50)

(51)

Colligative Properties

• Changes in

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



Boiling point elevation



Melting point depression



O

ti

(52)

Vapor Pressure

Due to both temperature

effects and energy transfers

f lli i l l

from collisions, molecules on the surface of a liquid are able to gain sufficient g kinetic energy to escape into the atmosphere

(53)

Vapor Pressure

• At any temperature, some molecules in a liquid have

enough energy to escape.

f f

• As the temperature rises, the fraction of molecules

(54)

Vapor Pressure

If the container is open to the atmosphere, the molecules simply

escape. This process is escape. This process is called evaporation.

As molecules escape

f h f h

from the surface, they take energy with them resulting in a cooling g g effect on the liquid.

(55)

Vapor Pressure

A desert water bag (left)g ( ) A desert canteen (center) An Army canteen (right)

(56)

Vapor Pressure

• Vapor pressure increases with increases with temperature.

• When the vapor

pressure of a liquid pressure of a liquid equals the

atmospheric

pressure the liquid pressure, the liquid boils.

• The normal boiling

f

point of a liquid is the temperature at which its vapor

(57)

Vapor Pressure

If the container is closed to the atmosphere, as

l l th li id th

more molecules escape the liquid, the pressure they exert increases.

(58)

Vapor Pressure

In a sealed container, eventually, the air space in the container becomes saturated with vapor molecules.

Th li id d h t t f d i

The liquid and vapor reach a state of dynamic

equilibrium: as liquid molecules evaporate, vapor molecules condense at the same rate. This is called the vapor pressure equilibrium

(59)

Vapor pressure of water at various temperatures

Temperature Pressure Temperature Pressure Temperature Pressure Temperature Pressure

°CC mm Hgmm Hg °CC mm Hgmm Hg °CC mm Hgmm Hg °CC mm Hgmm Hg 0 4.6 26 25.2 52 102.1 78 327.3 1 4.9 27 26.7 53 107.2 79 341.0 2 5.3 28 28.3 54 112.5 80 355.1 3 5.7 29 30.0 55 118.0 81 369.7 4 6.1 30 31.8 56 123.8 82 384.9 5 6.5 31 33.7 57 129.8 83 400.6 6 7.0 32 35.7 58 136.1 84 416.8 7 7.5 33 37.7 59 142.6 85 433.6 8 8.0 34 39.9 60 149.4 86 450.9 9 8.6 35 42.2 61 156.4 87 468.7 10 9.2 36 44.6 62 163.8 88 487.1 11 9.8 37 47.1 63 171.4 89 506.1 12 10.5 38 49.7 64 179.3 90 525.8 13 11 2 39 52 4 65 187 5 91 546 0 13 11.2 39 52.4 65 187.5 91 546.0 14 12.0 40 55.3 66 196.1 92 567.0 15 12.8 41 58.3 67 205.0 93 588.6 16 13.6 42 61.5 68 214.2 94 610.9 17 14.5 43 64.8 69 223.7 95 633.9 18 15 5 44 68 3 70 233 7 96 657 6 18 15.5 44 68.3 70 233.7 96 657.6 19 16.5 45 71.9 71 243.9 97 682.1 20 17.5 46 75.7 72 254.6 98 707.3 21 18.7 47 79.6 73 265.7 99 733.2 22 19.8 48 83.7 74 277.2 100 760 23 21.1 49 88.0 75 289.1 24 22.4 50 92.5 76 301.4 25 23.8 51 97.2 77 314.1

(60)

Vapor Pressure

Because of solute-solvent

intermolecular attraction

intermolecular attraction,

higher concentrations of

nonvolatile solutes make

it harder for solvent to

escape to the vapor

phase.

Therefore, the vapor

pressure of a solution is

lower than that of the

pure solvent

pure solvent.

(61)

Vapor Pressure

There are less molecules of the solvent in

the vapor phase above the solution

(62)

Raoult’s Law

The VP of H

₂₂

O (or the solvent) over a

solution depends on the number of H

₂₂

O

solution depends on the number of H

₂₂

O

molecules per solute molecule.

P

is

is proportional to

proportional to

X

P

_solvent_solvent

is

is proportional to

proportional to

X

_solvent_solvent

P

_solvent

= X

_solvent

P



_solvent

where

• X_solvent is the mole fraction of the solvent

• P_solvent is the normal vapor pressure of the solvent at that temperature

This eq ation is kno n as

Rao lt’s

Rao lt’s La

La

This equation is known as

(63)

Raoult’s Law

An ideal solution is one that obeys

An ideal solution is one that obeys Raoult’sRaoult’s law.law.

P_A_A = X_A_AP_A_A

•• Because mole fraction of solvent, Because mole fraction of solvent, XX_A_A, is always less , is always less

than 1, then

than 1, then PP_A_A is always less thanis always less than PPoo_A_A..

than 1, then

than 1, then PP_A_A is always less than is always less than PP _A_A..

•• The vapor pressure of solvent over a solution is The vapor pressure of solvent over a solution is

always

always LOWERED!

always

(64)

V

Vapor

Pressure

Lowering

(65)

Boiling Point Elevation and

F

i

P i t D

i

Freezing Point Depression

Nonvolatile solute-solvent interactions interactions cause solutions to have higher boiling points and lower freezing points g p than the pure solvent.

(66)

Boiling Point Elevation

The change in boiling point is proportional to

the molality of the solution:

T

_b

= K

_b

 m

where:

K_b is the molal boiling point elevation constant, a

property of the solvent.

(67)

Freezing Point Depression

• The change in freezing point can be found

similarly:

y

T

_f

= K

_f

 m

Where:

K_f is the molal freezing point depression constant of the

solvent.

(68)

Boiling Point Elevation and

F

i

P i t D

i

Freezing Point Depression

Note that in both

equations,

T does

T

b

= K

b

 m

not depend on

what the solute is

,

but only on

how

but only on

how

many particles

are

dissolved.

T

_f

= K

_f

 m

(69)

Colligative Properties of

El t l t

Electrolytes

Since colligative properties depend on the number of particles dissolved, solutions of electrolytes

(which dissociate in solution) should show greater changes than those of nonelectrolytes.

(70)

Electrolytes

Strong electrolyte: a compound that dissociates completely to ions in an aqueous solution.

Compound Dissociates to No. of ions per formula unit NaCl Na+ _{and Cl}- ₂ NaCl Na+ _{and Cl} ₂ CaCl₂ Ca2+ _{and 2 Cl}- ₃ K₂SO₄ 2 K+ _{and SO} 42- 3

Ionic substances dissociate into the ions and polyatomic ions

d i iti th h i l f l f th d

Mg₃(PO₄)₂ 3 Mg2+ _{and 2 PO}

43- 5

used in writing the chemical formulas of the compounds

Weak electrolyte: a compound that only partially dissociates to ions in an aqueous solution.q

An example is acetic acid, HC₂H₃O₂, which exists as HC₂H₃O₂ molecules, H+ _{and C}

(71)

Colligative Properties of

El t l t

Electrolytes

However,

a 1 M solution of NaCl does not

However,

a 1 M solution of NaCl does not

show twice the change in freezing point

that

a 1 M solution of methanol does.

(72)

The van’t Hoff Factor

The van t Hoff Factor

One mole of NaCl

in water does not

really give rise to

two moles of ions.

(73)

The van’t Hoff Factor

The 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.

(74)

The van’t Hoff Factor

The van t Hoff Factor

• Reassociation is more likely at higher

t ti

concentration.

• Therefore, the number of particles present is

t ti

d

t

concentration dependent.

(75)

The van’t Hoff Factor

The van t Hoff Factor

We modify the previous equations by

lti l i

b th

’t H ff f t

i

multiplying by the

van’t Hoff factor, i

(76)

Osmosis

• Some substances form

semipermeable

b

ll

i

ll

membranes

, allowing some smaller

particles to pass through, but blocking

other larger particles

other larger particles.

• In biological systems, most

semipermeable membranes allow water to

pass through, but solutes are not free to

do so.

(77)

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 solute concentration) to the area of lower solvent

(78)

Osmosis

The semipermeable membrane allows only the movement of The semipermeable membrane allows only the movement of solvent molecules.

Solvent molecules move from pure solvent to solution in an attempt to make both have the same concentration of solute attempt to make both have the same concentration of solute.

(79)

(80)

Osmosis

Dissolving the Egg in pure Egg in corn

Dissolving the shell in vinegar Egg in corn syrup Egg in pure water

(81)

Osmotic Pressure

• The pressure required to stop

osmosis known as

osmotic

osmosis, known as

osmotic

pressure

,



, is

Osmotic

n

V



=

( )

RT = MRT

Osmotic pressure

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 concentrations are the same), the solutions are isotonic.

(82)

Osmosis

••

Osmosis of solvent

from one solution to

th

ti

th

ti

another can continue

until the solutions are

ISOTONIC

—

— they

they

ISOTONIC

—

— they

they

have the same

concentration.

(83)

Normal red

blood cells

(84)

Osmosis in Blood Cells

• If the solute

t ti

concentration

outside the cell is

greater than that

inside the cell, the

solution is

hypertonic

.

• Water will flow out of

the cell, and

(85)

Osmosis in Cells

• If the solute

concentration

outside the cell is

less than that inside

the cell, the solution

is

hypotonic

is

hypotonic

.

• Water will flow into

the cell, and

(86)

Dialysis

Dialysis:

the separation

of larger molecules,

dissolved substances,

or colloidal particles

from smaller molecules,

substances or colloidal

substances, or colloidal

particles by a

semipermeable

p

membrane.

(87)

Water Treatment

(88)

Reverse Osmosis

W t

D

li

ti

Water Desalination

(89)

Reverse Osmosis

(90)

(91)

Colloids

• In true solutions, the maximum diameter

of a solute particle is about 1 nm.

• Colloid:

a solution in which the solute

particle diameter is between 1nm and 1000

p

nm.

• Colloid particles have very large surface

areas, which accounts for these two

characteristics of colloidal systems;

– they scatter light and, therefore, appear turbid,

cloudy, or milky.

– they form stable dispersions; that is they do – they form stable dispersions; that is, they do

(92)

Types of Colloids

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

(93)

Colloids

• John Tyndall (1820-1893)

• Tyndall effect:

Tyndall effect:

a characteristic of

colloids in which light passing

through the colloid is scattered

g

(i.e., reflected off of colloidal

particles).

Examples of colloids

that exhibit the Tyndall y

effect are smoke, serum, and fog.

(94)

Colloids

• Tyndall effect

(95)

Why is the sky blue?

Normal sky color

(96)

Colloids

• Robert Brown (1773-1858)

In 1827 the English botanist

• In 1827 the English botanist

Robert Brown noticed that

pollen grains suspended in

water jiggled about under

the lens of the microscope

the lens of the microscope,

following a zigzag path.

• Brownian motion:

the

random motion of

colloid-size particles.

(97)

Colloids

• Examples of Brownian

motion are the motion of dust particles in the air; what we see are the dust particles due p to scattered light.

(98)

Colloids

• Why do colloidal particles remain in solution

despite all the collisions due to Brownian

motion?

– Most colloidal particles carry a large solvation

layer; if the solvent is water as in the case of layer; if the solvent is water, as in the case of

protein molecules in the blood, the large number of surrounding water molecules prevents colloidal

molecules from touching and sticking together.

– Because of their large surface area, colloidal

particles acquire charges from solution; for particles acquire charges from solution; for

example, they all may become negatively charged. When a charged colloidal particle encounters