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Transcript
Chapter 13:
Properties of Solutions
Created By:
Nick Brazones, Brian Fike, Maya Hairston, Brian Kuttler, Alex Reardon, Chase
Schilling, Moses Suh
13.1 The Solution Process
•
•
A solution is formed when one substance
disperses uniformly throughout another.
With the exception of gas mixtures, all
solutions involve substances in a condensed
state.
13.1 Intermolecular Forces
•
•
Molecules or ions of substances in a liquid
or solid state experience intermolecular
attractive forces.
A major factor determining whether a
solution forms are the strengths of
intermolecular forces between and among
the solute and solvent particles.
13.1 Dissolution of an Ionic Solid in Water
•
NaCl dissolves readily in water because the
attractive interactions between the ions
and the polar H2O molecules overcome the
lattice energy of NaCl.
13.1 Dissolution of an Ionic Solid in Water
•
•
•
Water molecules orient themselves to the
surface of the NaCl crystals.
The positive end of the water dipole is
oriented toward the Cl- ions, and the
negative end of the water dipole is oriented
toward the Na+ ions.
The ion-dipole attraction pulls the ions
from their positions.
13.1 Dissolution of an Ionic Solid in Water
•
•
•
Once separated from the crystal,the Na+
and Cl- ions are surrounded by water
molecules.
Interactions like this, between solute and
solvent molecules, are known as solvation.
When the solvent is water, the interactions
are also referred to as hydration.
catalog.flatworldknowledge.com
13.1 Energy Changes and Solution Formation
The overall enthalpy change in forming a solution,
ΔHsoln, is the sum of three terms:
ΔHsoln = ΔH1 + ΔH2 + ΔH3
ΔH1: Separation of solute molecules (endothermic)
ΔH2: Separation of solvent molecules (endothermic)
ΔH3: Formation of solute-solvent interactions
(exothermic)
Remember: breaking bonds takes energy
http://wps.prenhall.com/wps/media/objects/3312/3391718/blb1301.html
Enthalpy Changes
http://wps.prenhall.com/wps/media/objects/3312/3391718/blb1301.html
13.1 Energy Changes and Solution Formation
•
Ionic solutes like NaCl do not dissolve in
non-polar solvents such as Gasoline,
because the non-polar hydrocarbon
molecules would experience only weak
attractive interactions with the ions.
13.1 Energy Changes and Solution Formation
•
Polar liquids, such as water, do not form
solutions with non-polar liquids like
gasoline, because the water molecules
experience strong hydrogen-bonding
interactions.
Solution Formation, Spontaneity, and Disorder
If you let go of a book, it falls to the floor because of gravity.
At its initial height, it has a higher potential energy than
when it is on the floor. Unless it is restrained, the book falls;
and as it does, potential energy is converted into kinetic
energy. When the book strikes the floor, the kinetic energy
is converted largely into heat energy, which is dispersed
throughout the surroundings. The book has lost energy to its
surroundings in the process. This fact leads us to the first
basic principle identifying spontaneous processes and the
direction they take.
Solution Formation, Spontaneity, and Disorder
First Basic Principle:
•
Processes in which the energy content of
the system decreases tend to occur
spontaneously.
Solution Formation, Spontaneity, and Disorder
Suppose we could suddenly remove a barrier that separates
500 mL of CCl4 from 500 mL of C6H14. Before the barrier is
removed, each liquid occupies a volume of 500 mL. When
equilibrium has been established after the barrier has been
removed, the two liquids together occupy a volume of 1000
mL. Formation of a homogenous solution has increased the
degree of dispersal, or randomness (entropy), because the
molecules of each substance are now mixed and distributed
in a volume twice as large as that which they individually
occupied before mixing.
Solution Formation, Spontaneity, and Disorder
Second Basic Principle:
•
Processes occurring at a constant
temperature in which the randomness or
dispersal in space (entropy) of the system
increases tend to occur spontaneously.
Solution Formation, Spontaneity, and Disorder
Third Basic Principle:
•
In most cases, formation of solutions is
favored by the increase in entropy that
accompanies mixing.
Solution Formation, Spontaneity, and Disorder
•
Consequently, a solution will form unless
solute-solute or solvent-solvent interactions
are too strong relative to the solute-solvent
interaction.
Solution Formation and Chemical Reactions
•
•
Be careful to distinguish the physical
process of solution formation from chemical
reactions that lead to a solution.
For example,
Ni(s) + 2 HCl(aq) → NiCl2(aq) + H2(g)
Solution Formation and Chemical Reactions
Ni(s) + 2 HCl(aq) → NiCl2(aq) + H2(g)
The chemical form of the substance
being dissolved is changed from Ni to NiCl2;
if the solution were evaporated, NiCl2, not
Ni would be recovered
13.2 Saturated Solutions and Solubility
•
•
•
As a solid solute dissolves, the
concentration of the solute in
the solution will increase.
If the concentration of the
solute gets high enough, the
particles may collide back
together.
This process is known as
crystallization.
catalog.flatworldknowledge.com
13.2 Saturated Solutions and Solubility
•
•
•
www.prevor.com
Solute + Solvent
Solution
The forward reaction is when
the solvent is dissolving.
The reverse reaction is when
the solute is crystallizing
back together.
13.3 Factors Affecting Solubility
http://www.engineeringarchives.com/les_matsci_unlimitedandlimitedsolubility.html
13.3 Polarity: Like Dissolves Like
•
•
Non-Polar substances tend to dissolve in
non-polar liquids.
Polar and ionic substances tend to dissolve
in polar liquids.
13.3 Polarity: Like Dissolves Like
•
•
Miscible: Pairs of liquids
that dissolve completely in
each other
Ex: Water + Acetone
Immiscible: Paris of liquids
that do not dissolve
significantly in each other
Ex: Water + Oil
http://library.thinkquest.org/CR0215471/oil_spills.htm
•
•
13.3 Polarity: Why Water is the
Universal Solvent
Water is polar which allows it to dissolve
polar and ionic solutions, which are the
majority of solutions.
Water has the ability to make hydrogen
bonds.
http://en.wikipedia.org/wiki/Hydrogen_bond
13.3 Polarity: Example Problems
Miscible or Immiscible?
Water + Sodium Fluoride
Vegetable Oil + Magnesium Chloride
•
•
13.3 Pressure: Trends
http://unit4lewiswain.blogspot.com/
•
•
The solubility of a
gas increases as
pressure of the gas
increases.
Does not affect the
solubility of solid or
solutions
13.3 Pressure: Henry’s Law
•
•
•
•
The relationship between pressure and
solubility of a gas is expressed as:
Cg = kPg
Cg is the solubility of the gas (molarity).
k is a constant that changes with the units
of volume and temperature (mol/ L*atm).
Pg is the partial pressure of the gas over the
solution.
13.3 Pressure: Example Problem
Given: At 25 C, .78 atm, 5.3*10-4M,
Pg is .05 atm
? Cg
•
•
Cg = kPg
Cg = (5.3*10-4 mol/L)/.78 atm (.05 atm)
Pressure: Example Problem Answer
3.4 x 10-5 M
13.3 Temperature: Trends
•
•
As temperature increases, the solubility of
most solids in water increases.
Exceptions do apply* Ce2(SO4)3
As temperature increases, the solubility of
gases in water decrease.
Ex: Boiling water releases gas molecules
13.3 Temperature: Solubility Chart
http://www.kentchemistry.com/links/Kinetics/SolubilityCurves.htm
http://www.barrymitzman.com/barry_mitzman/2012/01/mms-go-to-the-super-bowl.html
13.4 Ways of Expressing Concentration
Maya and Moses
13.4 Mass Percentage
● “soln” represents solution
13.4 Mass Percentage Problem
A solution is made by dissolving 13.5 g of glucose
(C6H12O6) in .100 kg of water. What is the mass
percentage of solute in this solution?
13.4 Part Per Million (ppm) and Part
Per Billion (ppb)
•
•
•
•
•
Diluted
1 ppm =106 grams = 1 milligram (mg)
Parts per billion (ppb)
1 ppb =109 grams = 1 microgram (µg)
Extremely diluted
13.4 Ppm And Ppb Problems
A 2.5 g sample of groundwater was found to contain 5.4 µg of
Zn2+. What is the concentration of Zn2+ in parts per million?
What is the concentration of Zn2+ in parts per billion?
13.4 Mole Fraction
● Symbol X is commonly used for mole
fraction, with a subscript to indicate
the component of interest.
● Example: XNACL
● Mole fraction of Sodium chloride
(Table Salt)
13.4 Molarity and. Molality
● Molarity (M):
● Molarity changes with temperature because the expansion or
contraction of the solution changes its volume.
•
•
Molality (m):
Molality does not vary with temperature because mass does not vary with temperature.
Thus, molality is often the concentration unit of choice when a solution is to be used
over a range of temperatures.
13.4 Molarity Problem
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.
13.4 Mole Fraction and Molality Problem
An aqueous solution of hydrochloric acid contains 36% HCL by
mass. Calculate the mole fraction of HCL in the solution.
Calculate the molality of HCL in the solution.
13.4
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13.5 Colligative Properties
•
Physical properties of solutions that depend
on the quantity but not the kind of solution
particles
Vapor pressure
o Boiling point
o Freezing point
o Osmosis
o
13.5 Lowering the Vapor Pressure
•
•
•
The pressure exerted by the
vapor
o A substance with no measurable vapor
pressure is nonvolatile whereas one
that has a vapor pressure is considered
volatile.
Adding a nonvolatile solute to a
solvent always lowers the vapor
pressure
o Expressed by Raoult’s Law
An ideal solution obeys Raoult’s Law
http://wps.prenhall.com/wps/media/objects/3312/3391718/blb1305.html
13.5 Raoult’s Law
•
PA=XAP◦A
o PA -Partial pressure exerted by solvent
vapor
o XA -product of the mole fraction of the
solvent in the solution
o P◦A -vapor pressure of the pure solvent
Example 1
•
Calculate the vapor pressure for a solution prepared
by dissolving 185.0g sucrose (mol. wt.=342.3) in 643.5
cm3 of water. the density of water is 0.9971g/cm3 and
the vapor pressure is 23.76 torr.
Example 1
First determine the number of moles of
sucrose:
Example 1
Next determine the moles of water:
Example 1
The mole fraction of water in the solution is:
Example 1
Then:
The vapor pressure has been lowered by
0.30 torr.
13.5 Boiling-Point Elevation
•
•
The boiling point of a solution is higher than
that of a pure liquid.
ΔTb=Kbm
o ΔTb- boiling point of pure solvent
o Kb- molal boiling-point-elevation constant
o m- molality
13.5 Freezing-Point Depression
•
•
The freezing point of the solution is lower
than that of the pure liquid.
ΔTf=Kfm
o ΔTf - freezing point of pure solvent
o Kf- molal freezing-point-depression
constant
o m- molality
Example 2
•
Calculate the boiling point and freezing
point of a 25.0 mass % solution ethylene
glycol (C2H6O2) in water.
Example 2
First calculate the molality of the solution:
Example 2
Calculate the changes in freezing and boiling point:
Example 2
So the boiling and freezing points are:
(normal bp of solvent) + ΔTb = 102.7◦C
•
•
(normal fp of solvent) - ΔTf = -10.0◦C
13.5 Osmosis
•
•
Some materials are semipermeable; they allow some
molecules to pass through but not others.
Osmosis- the net movement of solvent molecules from
the less concentrated solution to the more
concentrated solution
13.5 Osmotic Pressure
•
•
Osmotic pressure (π) is the pressure
required to prevent osmosis
Follows a law similar to the ideal-gas law
o
o
πV=nRT
π=MRT where M is the molarity of the solution
13.5 Osmotic Pressure
•
•
If two solutions have the same osmotic
pressure, no osmosis will occur and the
solutions are isotonic.
If one solution has a lower osmotic
pressure, it is hypotonic with respect to the
more concentrate, or hypertonic, solution.
Example 3
•
The average osmotic pressure of blood is
7.7 atm at 25◦C. What concentration of
glucose (C6H12O6) will be isotonic with
blood?
Example 3
•
•
•
π=MRT
M=π/RT
7.7atm/(0.0821)(298 K) = 0.31 M
13.5 Determining Molar Mass
•
•
The colligative properties of solutions can
be used to determine molar mass.
Any of the four colligative properties can be
used.
13.6 Colloids
•
•
•
Colloids are the intermediate types of
dispersions between solutions and
heterogenous mixtures.
The size of a colloid is one that is larger
than a molecule, but not large enough for
the components to be separated by gravity.
They can be found in any of the three
natural states of matter.
Examples of Colloids
•
•
Some examples of colloids are fog, smoke,
whipped cream, milk, paint, marshmallow,
butter, lotion, and ruby glass.
Milk is just a mixture of
fat droplets suspended in
water.
http://www.absoluteastronomy.com/topics/Colloid
13.6 Tyndall Effect
•
•
The diameter of a
colloid particle can
range between 5nm to
1000 nm.
Despite their size, they
are still large enough to
scatter light.
http://silver-lightning.com/tyndall/
13.6 Hydrophilic vs. Hydrophobic
•
•
The most important colloids are ones where
the dispersing medium is water.
The colloids are either hydrophilic or
hydrophobic.
13.6 Hydrophilic
•
•
•
A hydrophilic colloid is similar to large
molecules in the human body.
The hydrophobic groups are in the inside of
the molecule, while the hydrophilic groups
are on the outside of the molecule, keeping
it suspended in water.
They usually contain either oxygen or
nitrogen and carry a charge.
13.6 Hydrophobic
•
•
A hydrophobic colloid has to be stabalized,
because if not, the molecules will separate
in the water.
They can be stabilized by adsorption of ions
on their surface, or the adhesion of the ions
to its surface.
13.6 Hydrophobic Cont.
•
•
The ions on the outside of the
molecule are hydrophilic, so
they interact with the water
and keep it suspended.
The repulsion between the
ions and the colloid particles
keep the molecules from
colliding and getting larger.
http://www.life-treasure.com/technology.html
Example of Suspension
http://www.zazzle.com/chemists_are_frequently_in_suspension_humor_card-137220125589000317
Other Cases
•
•
Hydrophobic colloids can also be stabilized with
hydrophilic groups on their surfaces.
For example:
o Oil is a hydrophobic substance, so when in water, it
forms an oil slick.
o But in the presence of sodium stearate, a substance
with one hydrophilic end and one hydrophobic end, the
oil is stabilized.
o The hydrophobic end of the stearate interacts with the
oil, while the hydrophilic end interacts with the water,
suspending the molecule in water.
13.6 Emulsion
•
Emulsion is when a suspension of one liquid is
formed in another liquid.
o Example: process that occurs in the small
intestine
o In digestion, bile is released, and one specific
component in it emulsifies the fats, which
helps the absorption of them through the
intestine wall.
o The component in bile is an emulsifying agent.
13.6 Removal of Colloids
•
•
•
Sometimes, colloids must be removed from
the dispersing medium.
They are too small to be removed by
filtration, so they must be removed by
coagulation.
After coagulation, the colloids will be large
enough to filter, or they can settle out of
the dispersing medium.
13.6 Coagulation
•
•
Coagulation is when the colloids collide to
create larger molecules.
It can occur by either heating the mixture
or adding electrolytes
13.6 Heating
•
•
When heated, the speed of the particles
increase, allowing them to collide.
After they collide, the particles stick
together and grow larger.
13.6 Adding Electrolytes
•
•
•
When an electrolyte is added, it removes
the charge on the surface of the particles.
The neutral charge eliminates the
electrostatic repulsion between the
particles.
With no repulsion, they are allowed to
collide and grow larger.
13.6 Semi Permeable Membranes
•
Another way to separate ions from a colloid
is by using semi permeable membranes.
o
•
Ions can pass through the membrane, but the
colloids cannot.
This separation is otherwise known as
dialysis, and is used to purify blood in the
kidney.