Solutions

Solutions

• A solution is formed when one substance disperses uniformly throughout another.

• Intermolecular forces function between solute particles and the solvent molecules that surround them.

• Solutions form when the attractive forces between solute and solvent are comparable in magnitude with those that exist between the solute particles themselves or between the solvent particles themselves.

• (see pgh 2 on pg 470 B&L)

• When separated ions are surrounded by water molecules it is called solvation.

• When the solvent is water, the interactions are known as hydration.

Occur in all phases• The solvent does the dissolving.

• The solute is dissolved.

• There are examples of all types of solutes dissolving in all types of solvents.

• We will focus on aqueous solutions.

Ways of Measuring Concentration

• Molarity

• Molality

• % mass• Normality (just read this section)

• Mole Fraction

Molarity

• Molarity = moles of solute Liters of solvent

• Units are moles

liter

• Use the symbol “M” for molarity

Molarity

• Try some molarity calculations

• Molality = moles of solute Kilograms of solvent

• Units are mole/Kg

• Molality is abbreviated “ m “

Molality

Molality

• Try some molality calculations

% Mass

• % mass = Mass of solute x 100 Mass of solution

Normality

• Normality - read but don’t focus on it.

Mole Fraction

• Mole fraction of component A

A = NA

NA + NB

Dilution

• M1 V1 = M2 V2

• “M” is molarity

• “V” is volume (in liters)

• The overall enthalpy change in forming a solution ∆Hsoln, is the sum of the following:

• ∆Hsoln = ∆H1 + ∆H2 + ∆H3

Energy of Making Solutions• Heat of solution ( Hsoln ) is the energy

change for making a solution.

• Most easily understood if broken into steps.

• 1.Break apart solvent

• 2.Break apart solute

• 3. Mixing solvent and solute

1. Break Apart Solvent• Have to overcome attractive forces, this

requires energy. H1 >0 (endo)

2. Break apart Solute.• Have to overcome attractive forces. H2

>0 (endo) the solvent molecules need to separate to make room for the solute

3. Mixing solvent and soluteH3 depends on what you are mixing.

• Molecules can attract each other H3 is large and negative (∆H < 0 , exo) .

• Molecules can’t attract H3 is small and negative.

• This explains the rule “Like dissolves Like”

Energy

Reactants

Solution

H1

H2

H3

Solvent

Solute and Solvent

• Size of H3 determines whether a solution will form

H3

Solution

Types of Solvents and solutes

• If Hsoln is small and positive, a solution will still form because of entropy.

• There are many more ways for them to become mixed than there is for them to stay separate.

• The formation of a solution can be either exo or endothermic.

• If the process is exo it will tend to proceed spontaneously.

• A solution will not form if ∆Hsoln is too endothermic.

• The solute/solvent interactions must be strong enough to make ∆H3 comparable in quantity with ∆H1 + ∆H2

• Read pgh 2 and 3 on pg 472 B&L

• Two factors are involved in processes that occur spontaneously:

* energy

* disorder

• Processes in which the energy content of the system decreases tend to occur spontaneously. (spontaneous processes tend to be exothermic)

Types of Solutions• Solid solution-

Alloys are the most common solid solutions containing 2 or more metals

• Liquid solution-

Miscible - 2 or more liquids that can mix in any amount

Immiscible - liquids that cannot mix in any proportions

• Aqueous solutions have water as the solvent

Saturated Solutions and Solubility

• Several types of solutions:

• Unsaturated

• Saturated

• Supersaturated

Concentrations of Solutions

• Unsaturated -- A solution that contains less than the maximum amount of solute that can be dissolved at that temperature.

• Saturated Solution -- A solution containing the maximum amount of solute that can be dissolved at that temperature.

Supersaturated Solutions?

• Supersaturated -- A solution that contains more solute than would normally dissolve at that temp. Unstable!

• How can a solution be supersaturated? – Well, how can we dissolve MORE solute?– Heat!– So, heat a solution, dissolve MORE solute,

then cool it CAREFULLY.

A formerly supersaturated solution -- a single crystal of the solute introduced will cause ALL of the excess solute to come out of solution suddenly!

http://www.chem.ufl.edu/~itl/2045/lectures/lec_i.html

Factors Affecting Solubility

• Temperature

• Pressure

• Surface Area

• Agitation

Factors Affecting Solubility

Pressure

• External pressure has no effect on the solubility of solids or liquids because solids and liquids are not appreciably compressed when pressure is increased.

• Gases can be compressed easily.

Pressure - Solubility

• Gas compression increases the frequency with which gas molecules hit the liquid phase and enter it, thereby increasing the solubility.

• The effect of pressure on the solubility of gases is expressed in Henry’s law:

solubilitygas = k Pgas

Henry’s law

solubilitygas = k Pgas

• Solubility is expressed as molarity.

• “k” is the Henry’s law constant

• Pgas is the partial pressure of the gas.

Pressure effects

• Changing the pressure doesn’t effect the amount of solid or liquid that dissolves

• They are incompressible.

• Pressure does effect gases.

Dissolving Gases• Pressure effects the

amount of gas that can dissolve in a liquid.

• The dissolved gas is at equilibrium with the gas above the liquid.

• The gas is at equilibrium with the dissolved gas in this solution.

• The equilibrium is dynamic.

• If you increase the pressure the gas molecules dissolve faster.

• The equilibrium is disturbed.

• The system reaches a new equilibrium with more gas dissolved.

• Henry’s Law.

P= kC

Pressure = constant x

Concentration of gas

Factors Affecting Solubility

• Temperature• As a general rule, increasing temperature

drives molecules toward the more random phase. Therefore, an increase in temperature usually increases the solubility of solids in a liquid and always decreases the solubility of gases in a liquid.

• Decreasing temperatures have the opposite effect.

Temperature Effects

• Increased temperature increases the rate at which a solid dissolves.

• We can’t predict whether it will increase the amount of solid that dissolves.

• We must read it from a graph of experimental data.

• The next slide shows Solubility Curves for several substances.

20 40 60 80 10

0

Gases are predictable

• As temperature increases, solubility decreases.

• Gas molecules can move fast enough to escape.

Colligative Properties

• Colligative properties depend only on the number of dissolved particles in solution and not on their identity.

• • The four Colligative Properties are:• Vapor pressure reduction• Boiling point elevation• Freezing point depression• Osmotic pressure

Vapor Pressure

• A liquid in a closed container will establish an equilibrium with its vapor.

• When that equilibrium is reached, the pressure exerted by the vapor is called the vapor pressure.

• A substance that has no measurable vapor pressure is called non-volatile.

• A substance that has a vapor pressure is called volatile.

Vapor Pressure of Solutions

• A non-volatile solvent lowers the vapor pressure of the solution.

• The molecules of the solventmust overcome the force of both the other solvent molecules and the solute molecules.

Vapor Pressure• http://www.unit5.org/christjs/Vapor_Pressure_Boiling_Point.htm

A                                BIn container A, the liquid is evaporating.  Some of the molecules have enough kinetic energy to escape (turn to a gas) by pushing against the pressure of the atmosphere.  Container B shows the flask is saturated.  When new molecules of liquid are vaporized, the gas cannot hold additional molecules, therefore some of the molecules condense back to liquid.

Raoult’s Law:• Psoln = solvent x Psolvent

• Psoln is the vapor pressure of the solution

solvent is the mole fraction of the solvent

• Psolvent is the vapor pressure of the pure solvent

• This applies only to an ideal solution where the solute doesn’t contribute to the vapor pressure.

Raoult’s Law

• An ideal gas obeys the ideal gas law, and an ideal solution obeys Raoult’s Law.

• Real solutions best approximate ideal behavior when the solute concentration is low and when the solute and solvent have similar molecular sizes and similar types of intermolecular attractions.

Aqueous Solution

Pure water

• Water has a higher vapor pressure than a solution

Aqueous Solution

Pure water

• Water evaporates faster from pure water than solution

• The water condenses faster in the solution so it should all end up there.

Aqueous Solution

Pure water

Colligative Properties• Because dissolved particles affect vapor

pressure - they affect phase changes.

• Colligative properties depend only on the number - not the kind of solute particles present

• Useful for determining molar mass

Vapor Pressure & Boiling Point

•  Vapor pressure is the pressure exerted by a liquid in equilibrium with its pure liquid phase at a given temperature.

• The vapor pressure of a liquid is dependent only upon the nature of the liquid and the temperature.

• Different liquids at any temperature have different vapor pressures.

• The vapor pressure of every liquid increases as the temperature is raised.

• The normal boiling point of any pure substance is the temperature at which the vapor pressure of that substance is equal to 1 atmosphere (760 mm Hg).

Boiling Point Elevation

• The normal boiling point of a liquid is the temperature at which the vapor pressure reaches 1 atm. Since solutions exhibit vapor pressure lowering, the temperature at which the vapor pressure reaches 1 atm will be elevated, a phenomenon known as boiling point elevation.

• http://www.nyu.edu/classes/tuckerman/honors.chem/lectures/lecture_13/node7.html

Boiling point Elevation

• Because a non-volatile solute lowers the vapor pressure it raises the boiling point.

• The equation is: T = Kbmsolute

T is the change in the boiling point

• Kb is a constant determined by the

solvent.

• msolute is the molality of the solute

Freezing point Depression

• Because a non-volatile solute lowers the vapor pressure of the solution it lowers the freezing point.

• The equation is: T = Kfmsolute

T is the change in the freezing point

• Kf is a constant determined by the solvent

• msolute is the molality of the solute

1 atm

Vapor Pressure of solution

Vapor Pressure of pure water

1 atm

Freezing and boiling points of water

1 atm

Freezing and boiling points of solution

1 atm

TfTb

Electrolytes in solution• Since colligative properties only depend

on the number of molecules.

• Ionic compounds should have a bigger effect.

• When they dissolve they dissociate.

• Individual Na and Cl ions fall apart.

• 1 mole of NaCl makes 2 moles of ions.

• 1mole Al(NO3)3 makes 4 moles ions.

• Electrolytes have a bigger impact on on melting and freezing points per mole because they make more pieces.

• Relationship is expressed using the van’t Hoff factor i

i = Moles of particles in solution

Moles of solute dissolved

• The expected value can be determined from the formula.

• The actual value is usually less because

• At any given instant some of the ions in solution will be paired.

• Ion pairing increases with concentration.

• i decreases with in concentration.

• We can change our formulas to

H = iKm

• Label your solutions, in the flasks and the ice cube trays.

• Final conclusion will be to compare the actual freezing point depression to the theoretical.

• Give reasons for any differences.

• Liquid-liquid solutions where both are volatile.

• Modify Raoult’s Law to

• Ptotal = PA + PB = APA0 + APB

0

• Ptotal = vapor pressure of mixture

• PA0= vapor pressure of pure A

• If this equation works then the solution is ideal.

• Solvent and solute are alike.

Ideal solutions

Deviations• If Solvent has a strong affinity for solute (H

bonding).• Lowers solvents ability to escape.• Lower vapor pressure than expected.• Negative deviation from Raoult’s law.Hsoln is large and negative exothermic.

• Endothermic Hsoln indicates positive deviation.

How soap is used to dissolve non-polar solutes in polar

solvents.• Vocab terms:

• Hydrophilic – water loving

• Hydrophobic – water fearing

Soap

P O-

CH3

CH2CH2

CH2CH2

CH2

CH2

CH2

O-

O-

Soap

• Hydrophobic non-polar end

P O-

CH3

CH2CH2

CH2CH2

CH2

CH2

CH2

O-

O-

Soap

• Hydrophilic polar end

P O-

CH3

CH2CH2

CH2CH2

CH2

CH2

CH2

O-

O-

P O-

CH3

CH2CH2

CH2CH2

CH2

CH2

CH2

O-

O-

_

• A drop of grease in water

• Grease is non-polar

• Water is polar

• Soap lets you dissolve the non-polar in the polar.

Hydrophobic ends dissolve in grease

Hydrophilic ends dissolve in water

• Water molecules can surround and dissolve grease.

• Helps get grease out of your way.

Structure and Solubility

• Water soluble molecules must have dipole moments (polar bonds).

• To be soluble in non-polar solvents the molecules must be non-polar.

• HW: Brown/LeMay Ch. 13

• (10-12, 22, 23, 25, 28, 30, 32, 36, 42, 43, 45a, 46a, 49, 53, 57)

• Thanks to Mr. Green