[MUSIC]
Lighthouse Scientific
Education presents a lecture
in the Solutions
and Water series
The topic; Solutions
and Solubility.
The lecture is designed so
that the student can pick
and choice to view specific
sub-topics of the lecture.
As with many of our lectures,
not all students will need to
be exposed to all of
the material covered.
Be aware of what
is expected of you.
Material in this lecture
relies on an understanding of
the previous lectures
Fundamental
Properties of Matter
Ionic Bonding and
Intermolecular Forces.
The first part of the
lecture starts with some basic
definitions and
categories of solutions.
There are quite a few of them.
Then a look at
solution formation
with a specific focus
on forces and energy.
The second part of the
lecture deals with solubility
and is limited to solubility
with liquid solvents.
Factors effecting solubility
will be explored as will
the solubility of
solids and liquids.
The solubility rules
of ionic compounds in
water are reviewed.
The final topic is the
solubility of gases.
Henry's law will be introduced.
First to the definitions.
Some of these definitions we
have come across
before but some are new.
A term we have seen
before is solution.
It is a homogeneous
mixture of 2 or more
pure substances. Within
the solution is the solute.
It is a substance
that is dissolved in
another substance.
The other important
component of the
solution is the solvent.
It is the substance that
dissolves the solute.
Together the solute and
the solvent are a solution.
Perhaps the most familiar
solution begins with the
solute of salt;
sodium chloride.
Equally familiar is
the solvent H2O, water.
Adding salt to water, with
perhaps a bit of stirring,
creates the
solution salt water.
Solution is off
abbreviated s*o*l*n.
The salt is dissolved or
dissociated in the water and
together they form a solution.
Is there a limit to how
much salt can be dissolved
in a set volume of water?
Yes, there is and that
limit is referred to as the
solubility of
the salt in water.
Specifically, solubility
is defined as the ability of
a substance (solute),
to dissolve in a solvent.
A solute capable of being
dissolved in a solvent
is said to be soluble.
Some solutes are not
soluble in a solvent they are
said to be insoluble.
Sodium chloride is soluble in
water. But, only to a point.
That point, is
called its saturation.
A saturated solution
contains as much solute as
the solvent can dissolve.
The concertation is at the
saturation point.
Additional solute will
not mix homogenously
in the solvent.
If our cartoon salt
water solution represents
a saturated solution,
then the addition of any
more sodium chloride will
only result in the additional
solute sinking to the
bottom of the vessel.
The solute would be
referred to as being
a precipitant or
having precipitated.
A precipitation, abbreviated
ppt., has the solute in
a different phase
from the solution.
Rain is considered a
precipitation because it is
in the liquid form of
water as opposed to the
vapor phase in the cloud.
Returning to the salt water
solution, if the amount of
solute is less than that
of the saturation point the
solution is considered
an unsaturated solution.
By definition that means
the solution contains less
than the maximum amount
of solute that the solvent
is capable of dissolving.
Slightly outside of
saturated/unsaturated
definitions sits
super saturation. That
is when more solute
is dissolved in solution
than normally allowed.
It is condition
dependent and unstable.
It is also a topic that doesn't
get much play at this level
of chemistry and we will
not mention it again.
Sometimes there is
confusion with the
term solution.
Often students think that
a solution has to be a liquid.
But that's not the case.
Referring back to the
definition of solution, it is
a homogeneous mixture of
2 or more pure substances.
Nowhere in there does it
specify the states of matter.
A solution can actually
be a gas, liquid or solid.
A table that catalogs
solution, solute and solvent
by state of matter
should highlight
the diversity of
the term solution.
Starting with a
solution that is a gas.
For a solution to be a gas
both the solvent and the
solute have to be gases.
An example is air. Nitrogen
is the most abundant form
of gas in our atmosphere
and can be considered
the solvent. A less
abundant form of gas
is oxygen which
is called a solute.
There are additional solutes
in the solution we call air.
For solutions that
are liquid there is a
liquid solvent
and liquid solute.
Rubbing alcohol is a
good example of this
type of solution.
It contains both water
and isopropyl alcohol.
The alcohol is usually in
greater quantity and is
considered the solvent.
An important term
associated with this type
of the solution is miscible.
It implies that the liquids
involved in the solution,
the solvent and solute,
will mix completely
together in all, or
nearly all, proportions.
Contrast that with immiscible.
These are liquids that
do not mix completely
and generally form
separate layers.
Oil and water is a good
example of immiscible liquids.
Think of how salad dressing
will separates into different
layers when left to sit.
There are other types of
solutions that are liquids
including those that are
composed of liquid solvent
and a gas solute.
Carbonated water or soda
is a good example. Yes,
some of the gas, as bubbles,
does leave this type of
solution but there is always gas
remaining dissolved
in the solvent.
The third type of liquid
solution has a liquid
solvent and a solid solute.
The salt water example used
earlier serves well here.
The final two types of
solutions are solids.
This is where the common
use of the term solution as
a liquid comes in conflict
with its use in chemistry.
There are solid solutions
that are composed of
solid solvent and solid solute.
Metallic compounds such
as brass or steel are this
kind of solution. Steel
can also be categorized as
a compound so there can
be overlap in definitions.
These solutions are also
referred to as alloys.
The definition of an alloy
is a substance composed
of 2 or more metals or
a metal and a nonmetal.
The last solution in the
table is a solid that has
a solvent that is solid and
the solute that is liquid.
These solutions are
not as common as the
other solutions. An example
is the dental amalgam that is
used to fill cavities in teeth.
While this is a busy table
the solutions that get the
most attention in
this lecture are
the liquid solutions with
solutes as gases or solids.
We will be returning
to these solutions after
finishing up with a
few more definitions.
These definitions represent
a broadening of the previous
definition that has a solution
as a homogeneous mixture
of 2 or more pure substances.
There are two
important kinds of
solutions that are not
homogeneous mixtures.
The fall under the heading
heterogeneous solutions.
The first one is a suspension.
That is a fluid containing
solid particles that will
eventually settle out.
Consider this cartoon
representation of the
solution that has particles,
colored as yellow dots,
that are distributed
throughout the solution.
After some period of time,
if undisturbed, the particles
will settle, or sink to
the bottom of the solution.
The particles are not
considered dissolved
in the solution
but rather mixed in.
Common suspensions include
Sand in water and
dust in the air.
Given sufficient time both
types of particles will
settle out of the solution.
Yes, they can be
redistributed in the solution
with some agitation.
The other type of
heterogeneous solution
is called a colloid.
This is a solution with small
particles of varying size,
on the order of 1 to 100
nanometers that are evenly
distributed throughout.
A murky appearance is often
associated with colloids.
Because the particle size is
so small it is very
difficult to filter them out.
Also, because of the small
size, the particles tend to
stay suspended in solution.
This cartoon representation
has particles with a
distribution of sizes.
Commonly found colloids
include fog.
Water droplets of different
sizes (liquid) evenly
distributed within the
air (a gas solvent).
Smoke, tiny particles of ash
(solid) distributed in the air.
Emulsions, this is where
small particles of one liquid
are evenly distributed
within another.
Milk and mayonnaise can
be considered emulsions
since it has small drops
of fat distributed within
a water base.
The same goes for
cheeses and jellies.
That concludes the definition
portion of the lecture.
Solution formation
can be approached from
a mechanical perspective
using the lens of energetics.
Favorable energetics
has solutes dissolving in
a solvent. Unfavorable
energetics describes
an insoluble solute.
The energetics of a solution
formation can be broken
down into three parts.
These parts are not a
description of the actual
formation of the solution but
rather a way to consider the
costs and returns of creating
a solution. It goes like this.
The first part is the cost
or energy of separating some
of the pure solvent
particles from each other.
Breaking intermolecular bonds.
Cartoon example has
a beaker of solvent in
the liquid form. Liquid
is used for convenience
and depicted as
spherical particles.
Particles in the liquid
are attracted to each other
(intermolecular bonding)
but can still move freely
past each other.
It will be necessary to
break some of those
interactions so that the
solute particles can
interact with the solvent.
The volume has been
increased in anticipation
of the solute particles.
Breaking intermolecular
interactions costs energy.
In the same vein there
is the energy cost of
separating solute particles.
Solute particles can also
be thought of as initially
existing in a beaker with
some sort of intermolecular
interactions occurring
between the particles.
When they are in the
solution some of those
intermolecular interactions
will need to be broken so
that new solvent-solute
interactions can be formed.
Energetically, breaking
intermolecular bonds
costs energy.
If this were the only two
steps then there would
never be solution formation
because the process
would only involve steps
that cost energy.
But sometimes there
is solution formation
so a third step is needed
and that is for solute
and solvent particle to
interact. These
interactions releases energy.
Not always a lot but matter
interacts with matter.
With these 3 energy terms
the overall energetics of
solution formation
can be examined.
There are two scenarios.
If the magnitude of the
combined costs of steps
1 and 2 is less than the
magnitude of the energy
released in step 3
the solute is soluble
in the solvent.
Solution formation
is favorable.
If, however, costs
of steps 1 and 2
is GREATER than the
energy released in step 3
then the solute is NOT
SOLUBLE, or insoluble,
in the solvent. It just
costs too much for what
is gotten by mixing.
This concludes the
solution parts of the
lecture. It is also
the starting spot for the
second major part of the
lecture; solubility. Breaking
down the process of solution
formation by energetic
steps is useful in evaluating
factors affecting solubility.
In this lecture we
will limit the discussion
to solvents that are liquids.
As mentioned earlier,
solubility is the ability
of a substance, a solute,
to dissolve in a solvent.
Another way of saying that
is the solute is capable of
being dissolved.
A substance is defined
as soluble if more than 0.1 g
of it will dissolves in 100 ml
of solvent. Solubility, and
factors effecting solubility,
can be viewed from
several perspectives.
Listed here is an overview of
factors or perspectives that
are covered in the solubility
part of the lecture.
We will spend some
time with each one.
One aspect of solubility
is the speed of dissolving
of the solute. There
are factors related
to the properties of the
solute and solvent and there
are mechanical ways of
speeding up the process.
And then there is the most
useful phrase in solubility;
'like dissolves like'.
The phrase reduces
favorable solute-solvent
interactions to polar-polar
or nonpolar-nonpolar.
Temperature also plays a role.
The temperature
dependence of solubility
for liquid solutes is generally
opposite of that for gases.
Each will be addressed
in its own section,
Pressure dependence
also varies between the
liquid and solid solutes
and the gas solute.
First to factors the effect
the speed of dissolving.
It should be noted that these
factors are generalizations
and are not absolute.
The biggest driver in the speed
of dissolving is the amount
of surface area of the solute.
Since the favorable
formation of a solution is
a function of the
solute-solvent interactions
it should not be too
surprising that the more
accessible surface area
of the solute will mean
more solute-solvent
interactions.
The more solvent molecules
that can intermolecular bond
with a solute particle
the greater the release
of energy. This
often translates
into faster dissolving.
With that in mind,
it is not uncommon for
a solute to be crushed as
a way to increase
surface area and speed
the process along.
Another helpful
technique is stirring.
It uses the same argument
in that it produces more
solute-solvent interactions
which is synonymous with
faster dissolving. Bottom line,
the more contact the
solute has with the solvent
the faster it dissolves.
The next concept with
solubility is perhaps the
most famous statement
in chemistry.
It is 'like dissolves like'.
And by like or dislike the
criteria is between polar
and nonpolar (apolar is also
a term used for nonpolar).
Specifically, polar solvents
readily dissolve polar solutes.
Polar solvents
like polar solutes.
And nonpolar solvents
dissolve nonpolar solutes.
Nonpolar solvents
like nonpolar solutes.
By extension, nonpolar
substances do not like or
dissolve in polar substances.
This phrase is faster and
easier to use than an analysis
using the three energetic
steps as listed above.
Still, within the refrain
'like dissolve likes' are
the two cost steps and the
one energy releasing step.
For example, a nonpolar
solute should not dissolve
in a polar solvent.
They are not 'likes'.
Energetically, the cost of
separating the particles
of the polar solvent
and the cost of separating the
nonpolar solute particles is
not sufficiently paid for
with the release of energy of
the polar-nonpolar
intermolecular interaction
of the solution.
The insolubility of
the nonpolar solute in the
polar solvent is the result of
costs exceeding returns.
What about solute molecules
that have both polar and
nonpolar characteristics?
These molecules are not as
easy to categories as
soluble or insoluble in
a particular solvent.
They contain both like
and dislike qualities.
A look at how well different
organic alcohols dissolve
in water will show that the
ratio of polar to nonpolar
character determines the
level of solubility. Consider
the smallest organic
alcohol methanol, CH3OH.
It has a nonpolar
methyl group, CH3, and
a polar hydroxyl group, OH.
The hydroxyl group is
the alcohol.
Organic alcohols will
dissolve, or mix, into
the polar solvent water to
the degree to which
the polar component, the
hydroxyl group, is of
a sufficient proportion
as compared to the nonpolar
component (as measured
here by the amount of
carbon-hydrogen bonds).
In the case for methanol the
hydroxyl group comprises a
sufficient proportion
of the molecule so that
it completely dissolves
in the polar solvent.
It will do so at
all concentrations.
It is infinitely soluble.
The two liquids are
said to be miscible.
But what happens with
alcohols that have
larger nonpolar components?
This can be shown by
increasing the number of
carbons and hydrogens in
the alcohol while keeping
the polar characteristic
constant (a single
hydroxyl group).
It should be noted that
C2H5OH is ethanol which
is the compound commonly
referred to as alcohol
and is the substance found
in beer and wine.
Descending the table has
the alcohols acquiring
proportionally more
nonpolar character.
Back to the ethanol,
it also has enough polar
characteristic to
completely dissolve in water
in all proportions.
Same thing with
propanol, C3H7OH.
It is not until butanol,
C4H9OH, that we come
across an alcohol
with sufficient nonpolar
characteristic to limit the
solubility of the
alcohol in water.
Here the solubility of butanol
is 8.0 g per 100 ml of water
at 25 degree C. It is
not completely miscible in
water but makes the
cut of as soluble.
The larger nonpolar
characteristic of the alcohol,
4 carbons bonded to
9 hydrogens, limits
its solubility. Also
note that a solubility
value has to be given at
a particular temperature.
More on this latter.
Returning to the table,
having even more carbons
further limits the solubility.
At some point the alcohol
no longer is considered
a 'like' with water.
The take home trend is that
the more carbon-hydrogens
a substance has, the
more nonpolar it is,
and the less soluble
it is in water.
The next solubility trend
under the 'likes dissolve likes'
banner is salts or
ionic compounds in water.
It is a topic that was
originally discussed in the
'Ionic Bonding' lecture
and further developed in the
'Intermolecular
Forces' lecture.
When dealing with the
solubility of an ion in water
the term dissociation
is the used.
Dissociation is more
than just dissolving.
For ionic compounds
dissociation is dissolving in a
solvent into its
constituent ions.
The term dissolving can
still be used with ionic
compounds but it
implies dissociation.
Solubility is the
degree to which the
compounds dissociate.
Another way to think of
the solubility of
an ionic compound
is that the more
soluble a compound is,
the larger the concentration
of dissociated ions is at
the saturation point.
More soluble means
more dissociation.
The example used in the
previous lecture was the
salt sodium chloride
dissolving or dissociating
into solvated ions.
As said, polar solvents
dissolve polar solutes.
Remember, that polar bonds
have partial charges.
Polar-polar interactions are
electrostatic interactions
of the partial charge.
It is not too big of a leap
to extend the electrostatic
interaction into full charges.
So we will expand the
description into polar solvents
dissolve charged solutes.
A closer look at the
dissociation shows that
the water orients
itself around the cation
so that the partial negative
charge on the oxygen is
pointed toward the positive
charge of the cation
and the partial positive
charge of the hydrogen is
oriented toward the
negative charge of the anion;
likes dissolving like.
As for the basic trends in
dissociation and solubility
of ionic compounds
the 3 steps of the energetic
process serve as a guide.
Specifically, we will consider
the effects on solubility of
increasing size of the charge
on the ions in a salt as well
as increasing physical size
of the ions. To
make matters a bit
simpler we will only
consider aqueous solutions.
That will allow us to
ignore energetic step 1 when
considering the trends since
separating the molecules of
the solvent water will, on a
first pass, be similar for the
dissociation of
all ionic compounds.
Starting with the trend
regarding increasing
size of charge.
What is relevant in
describing this trend is the
energetic cost of separating
the anions and the cations.
Since ionic bonding is a
matter of positive charge
attracted to negative charge
the size of those charges is
directly proportional to the
strength of the bond.
Larger charges form
stronger bonds and
stronger bonds take more
energy to separate.
The cost (step 2) of solution
formation goes up.
The trend concerning the
solubility or dissociation
of ionic compounds in water
as a function of ion charge
is that the larger the charges
in the compound the
lower the solubility.
Compounds with large
charges have a saturation
concentration that is
lower than compounds with
smaller charges . Again
that is a generalization.
As for the increasing size
or radius of the ions in an
ionic compound, the driving
force in solubility is
the energy release found
with the solute-solvent
interaction. As
stated earlier, larger
solute particles have more
surface area to participate
in intermolecular
bonding with the solvent.
More intermolecular
bonding means more
energy release. The
trend concerning the
solubility or dissociation
of ionic compounds in water
as a function of ion size is
that the larger the ions in
the compound the
higher the solubility.
Compounds with large ions
have a higher saturation
concentration
than compounds with
smaller ions. And that
is another generalization.
We can take this
basic understanding of
dissociation of ionic
compounds and review
the common solubility rules.
These rules are given
in the student's
textbook and are generally
used when writing
net ionic equations.
While there is some
variation in rules between
different textbooks they
mostly follow the logic
of this set of rules.
Not all students will
need to understand the basis
of the rules but we will
cover them anyway.
Rule 1: All alkali metal
salts are very soluble
The alkali metal are column
1 in the Periodic Table.
They form plus 1 cations.
That is the smallest full
positive charge. They are
soluble because they have
a small charge.
The alkali metals will start
our table. They will
be given the solubility
rating of very.
The only polyatomic positive
ion we come across in
our studies, ammonium
(NH4+), forms salts that
are also very soluble.
Ammonium is a big ion
and has a small charge.
Both of those features
promote solubility.
And then there are
the halogen ions.
Chloride, bromide,
iodide they all form very
soluble salts. The
smallest halogen fluoride
is considered a
relatively insoluble ion.
The Periodic Table shows
the halogens to be in column
17 or 7B. Halogens
form -1 anions.
That is the smallest
full negative charge.
There are some
exceptions to this rule.
Halogens that form
compounds with the
cations silver I,
lead II or mercury I
(usually founded as Hg2
with a +2 charge) will be
considered as insoluble.
As for polyatomic
anions, sulfates salts
(SO4-2) are very soluble
All except for salts with
lead II, mercury II, calcium,
barium or strontium. The
last 3 are cations from atoms
in column 2 of
the Periodic Table.
These alkaline earth
metals form +2 ions.
Many of the single
charged polyatomic anions
form soluble salts.
These are salts with nitrate,
chlorite and chlorate.
Also the larger acetate ion
forms soluble salts.
Hydroxides are a
bit more complicated.
Generally, they are insoluble
but here is an important tip
regarding a table
of solubility rules:
the rules above
override the rules below.
Therefore, according to rule 1:
hydroxide salts of the
alkali metal are also soluble.
These are added as exceptions.
There are a few more
hydroxide salts to consider.
Those alkaline earth metals
that form +2 ions, well they
form moderately or slightly
soluble salts with a
hydroxide anion.
The common -2 or -3
polyatomic anions like
carbonate, phosphate,
sulfate and chromate
are generally considered
insoluble. -1
polyatomic ions are
generally soluble and
larger -2 or -3 charged
polyatomic ions are
generally insoluble.
Except for salts with the
alkali metal cations of rule 1.
Top rules override
bottom rules.
Finally, sulfides, S -2,
are insoluble.
Sulfides are small and
have a moderate charge.
This doesn't bode well
with the solubility trends
covered early. And
like the rules covering
hydroxides, sulfide salts
with alkali metal cations and
some of the alkaline earth
metals cations are going to
be moderately or very soluble.
There are patterns in this table.
It is worth considering
the rules in context with
each other and the energetic
steps of solution formation.
The last 2 factor on our
list affecting solubility are
temperature and pressure.
Our coverage of these
factors will be broken down
by states of matter.
Temperature and pressure
will be covered together
for solids and liquids and
covered separately for gases.
Starting with
solids and liquids.
In all cases the solubility
of a particular solute-solvent
mixture is given at a
specific temperature.
It is usually noted alongside
the saturation concentration.
Solubility in liquid of most
solids and liquids solutes
increases with
higher temperatures.
This is because at higher
temperatures the particles
in the pure solvent and
the pure solute have more
thermal motion and therefore
find it easier to overcome
some of their intermolecular
bonds. That is, the energetic
costs of steps 1 and 2
go down with temperature.
This argument can be
expanded to include the
interaction in step 3 but
the drop in energy for steps 1
and 2 are the
driving force here.
As shown with this graph,
the solubility of compounds
like nitric acid and
potassium chloride increase
with increasing
temperature (x-axis).
More of these compounds
can be dissociated
at higher temperatures.
They have high
saturation concentrations.
However, as seen with the
cerium sulfate, not all
compounds follow this trend.
The solubility of
cerium sulfate actually
decreases with
increasing temperature.
Trends can be helpful in
understanding solubility as
a whole but they should
not be taken as absolutes.
Because the addition of a
modest amount of pressure
produces very little change
in the density of a solid or
liquid we can generalize that
pressure has almost no effect
on the solubility of
liquid or solid solutes in
a liquid solvent. So, for
liquid and solid solutes.
'increase temperatures
increase solubility'.
And for the most part
pressure does not play
a big role in solubility.
Gas as a solute
is another thing.
The solubility of a
gas in a liquid solvent
is a balance between:
gas particles dissolving
in the solvent.
We will call that
as being captured and
dissolved gas particles
escaping from the solvent.
We are familiar with
gas moving in and out
of a solvent in the form
of a carbonated beverage.
Carbon dioxide is a gas
found in the bubbles which
are escaping from the drink.
Before we discuss the effect
of temperature and pressure
on the solubility of a gas it
might be helpful to expand
on our understanding of
what it actually means
for a gas to be soluble
in a liquid solvent.
For that we will use
a cartoon container that has
a liquid solvent with gas
solute particles above and
within the solution.
A lid is include to
eliminate evaporation.
That is discussed in more
detail along with
vapor pressure in the
'Colligative
Properties' lecture.
Returning to the solvent,
it is colored dark blue here.
The molecules that
make up the solvent are
constantly moving around.
They have thermal
energy of motion and are
continuously making and
breaking intermolecular bonds.
Now consider
a gas particle above the
surface of the liquid.
It is freely moving
about as would be expected
for a gas particle.
Such a particle has
a very high energy of motion.
Much more than that of
the solvent molecules.
For the most part these
zipping particles do not
participate in
intermolecular bonding.
This is only slightly
untrue but we can argue no
interactions because
they are very short lived.
As such, when considering
the energetics of solution
formation the cost of
separating solute particles
(step 2) is effectively zero.
There are no costs here.
Returning to the motion
of our gas particle,
at some point the
gas particle will bang
into the solvent.
If it is going real fast it
will probably bounce off.
If it is one of the slower ones
it may actually interact
with the solvent.
and dissolve into the solvent.
The solvent has captured
the gas particle.
While in the solution the
gas particle will move around
like the solvent molecules
do because it has energy
of motion. Less
than it did when it
was free above the solution.
It is also continuously
making and breaking
intermolecular bonds with
the solvent. These
are competing forces.
Motion verses bonding.
Some fraction of gas
particles, however, will
have enough thermal energy,
motion, to overcome the
intermolecular bonding with
the solvent and escape
from the solution back
into the gas phase where
the whole process of capture
and escape begins again.
Solubility is the
balance between
capture and escape.
One could think gas solubility
as measure of the amount of
gas in solution (captured)
compared to the amount of
gas out of solution (escaped).
More soluble gases are
more captured gases.
With this understanding of
escape and capture we can
devise some basic trends
for the solubility of gas.
Consider the
component of capture.
Anything, be it condition or
substances, which reduce the
capture component of
the process will tend to
lower the solubility.
Since capture is what
gets gas into solution it
makes senses that limiting
this limits solubility.
Conversely, anything
that increases the capture
aspect tends to increase
the solubility. It's the reverse
of the previous argument.
Now anything the decrease
the capacity of the
gas to escape the solution
should lead to an increase
in solubility.
More gas molecules
will remain in solution.
And anything that promotes
gas escaping will have the
opposite effect and
drive down solubility.
We are now ready to see
what affect temperature
and pressure have on
solubility of a gas in a liquid.
First to temperature and
a quick question to make
a point. Consider
two cans of soda.
Two cokes if you will.
The first one comes straight out
of an ice chest full of
ice water and the second
has been sitting on a table.
Which one will fizz
more when opened?
That is the one with the
lower solubility and it
proves the trend with
solubility of gases.
The solubility of a gas (the
solute) in a liquid solvent
decreases with
higher temperatures.
That is the exact opposite of
the trend for liquid and solid
solutes in a liquid solvent.
The key to understanding
this behavior is that
increasing temperatures
means energy of motion,
higher energy of motion.
Consider the gas phase.
Higher energy of motion
is a faster gas particle.
The faster the particle moves
the faster it will be traveling
when it strike the
surface of the liquid.
That means it is less likely
that the solute-solvent
intermolecular interaction
at the surface will be strong
enough to capture
the gas particle.
Slower particles get
captured more readily than
fast ones do. Capture
of the gas particles
goes down with increasing
temperature and therefore
solubility is
expected to go down.
But that is only half of it.
There is still the
gas particle dissolved
in the solvent.
They too have higher energy
of motion (not as much as in
the gas phase) and
therefore a larger
proportion of these gas
particles will have sufficient
energy to break free of the
intermolecular interactions
of the solution.
Escape goes up.
There seems to be
agreement with both capture
and escape: the solubility
of the gas goes down with
increasing temperature.
That can be seen
experimentally by looking at
the solubility's of 4
noble gases in water with
increasing temperature (x-axis).
It is not always
a prominent decline
but it does not go up.
The effect of pressure on
solubility can be argued
in the same manner.
We start by declaring
that pressure has a large
effect on gas behavior
in general and that comes
out with its solubility.
Increasing pressure
increases the solubility of
a gas in a liquid solvent.
This follows the line of
reason that more pressure
means more gas particles
for some given area.
This comes right out of the
definition of pressure of a gas
In the gas phase, with more
gas particles zipping around
there will be more particles
that have the right energy to
be captured by the solvent.
It is basic statistics.
More particles means that
there are more fast particles
and more slow particles.
There are more particles
with the right energy.
Capture goes up and
so does solubility.
At least as far as the
gas phase is concerned.
Due to the higher capture
there are more dissolved
gas particles in the solvent.
They too have thermal
energy and move around.
Since there are more
dissolved gas particles,
there are more particles with
enough energy to escape.
Escape will go up which
suggests that solubility
should go down.
There appears to be
conflicting influence on the
solubility due to pressure.
The gas phase argues
pressure increases solubility
while the dissolved phase
argues the gas should
be less soluble. What gives?
When the pressure is
increased it means more gas
particle in the gas phase
and the dissolved phase but
there are many more
particles in the gas phase
so there are more captures
than escape resulting
in a net flow into
the solution and a net
increase in solubility.
When dealing with the
solubility of gases in liquids,
an increase of temperature
decreases solubility and
an increase in pressure
increases solubility.
The final component
in a discussion of solubility
of gases, and not all students
will need to know this, is
Henry's Law. Again,
consider gas particles
moving above and
within the solution.
The movement of particles
above the solution
is the pressure of the gas.
If there are other types of
gases above the liquid then
the pressure due to any one
particular type of gas is
called its partial pressure.
In 1803 William Henry
noted that, at constant
temperature, T, there was a
direct relationship between
the solubility, S, of a gas
and it partial pressure, P.
We call it Henry's Law.
The specific wording is
"the amount of a given gas
(the solute) that dissolves in
a given type and volume
of liquid (the solvent) is
directly proportional to the
partial pressure of that gas
in equilibrium
with that liquid."
The solubility of a gas is
proportional to the
pressure of that gas.
The proportionality
constant, kH,
has to be provided.
Since kH is an
experimentally determined
value a table is made
available that contains
gas specific constants.
Having access to the
constants means that
Henry's Law only has 2
unknowns (solubility and
pressure). Most Henry's Law
questions give a value for
one of the unknowns
and the student solves
for the other.
That should be doable
at this level of chemistry.
And that completes the
material for this lecture.
Recapping the lecture.
It had two parts;
solution and solubility.
For solutions, first and
foremost, a solution is
homogeneous mixture of 2
or more pure substances.
The solute is the substance
dissolved in another
substance. The solvent
is the substance
that dissolves the solute
Solubility is the ability
of a substance (the solute),
to dissolve in a solvent.
We can also say that a
soluble substance is
capable of being dissolved.
A saturated solution has as
much solute as the solvent
can dissolve. The addition
of any extra solute will
precipitates or
drop out of solution.
It will not mix homogenously.
An unsaturated solution
has less than the maximum
amount of solute that
the solvent can dissolve.
Under certain circumstances
a solution can be considered
supersaturated. It
will have more solute
in solution than
is normally allowed.
It is usually unstable.
The term miscible is used for
liquids that can mix in all
(or nearly all) proportions.
Small alcohols are
miscible in water.
Immiscible is for liquids
that do not mix completely
or form separate layers.
Think oil and water.
Additional definitions
that fall outside of
'homogeneous mixture'
include 'suspension'.
These are fluids that contain
solid particles that will
eventually settle out.
And then there is 'colloid'.
It is a solution with small
particles of varying size
(from 1-100 nm) that are
evenly distributed
throughout the solution.
Tell-tale signs are a
murky solution that does
not filter well.
We spent some time
discussing the energetics of
solution formation.
It was broken down into
3 steps or considerations.
The first two steps are
the energetics of separating
the particles before mixing:
solvent and solute particles.
These steps cost energy.
The third is the
energy releasing step.
It is the energetics
of solute and solvent
intermolecular interactions.
Now, if the sum of
steps 1 and 2 is greater than
the energy release of step 3
then the solute is considered
insoluble in the solvent.
However, if the energy
release of step 3 is greater
than the sum of steps 1 and 2
then the solute is considered
as soluble in the solvent.
Continuing on to the
second part of the lecture;
factors affecting
solubility in a liquid.
We concentrated on 3 of them.
The first was the popular
refrain 'likes dissolves like'
That is, nonpolar solvents
are good at dissolving
nonpolar solutes
and polar solvents are good
at dissolving polar solutes.
Charged particles dissolve
in polar solvents and we
saw a couple trends
for the solubility of
ionic compound in water.
The larger the charge on an
ion the less soluble
the salt tends to be.
The cost of separating
the solute is higher.
As the size of the
ion increase so does
the solubility. Bigger
ions can participate
in more intermolecular bonds.
These trends were used to
sum up the solubility rules
for ionic compounds in water.
Solubility rules are used
with net ionic equations.
Another trend we looked at
under the 'like dissolves
like' banner was the
solubility of alcohols
in water. Alcohols
have both polar and
nonpolar regions.
The nonpolar region
being the
carbon-hydrogen bonds.
The trend is that solubility
decreases with increasing
amount of the nonpolar
part of the molecule:
the C-H bonds.
Temperature plays a
big role in solubility.
In fact the solubility value
for a solute is given at
a specific temperature.
We saw the trend for the
solubility's of most solids
and liquids increases
with higher temperatures.
That is in contrast to
the solubility of gases.
They decrease at
higher temperatures.
We used the perspective
of gas particles being
captured by, or escaping
from, the solution as
a means of explaining the
temperature and pressure
dependence on solubility.
Factors that inhibit
capture decrease solubility
while factors that promote
capture increase solubility.
Factors that inhibit escape
of gas particles increase
solubility and factors
that promote escape
reduce solubility.
As for pressure, it has little
effect on solubility of solids
and liquid solutes.
However, increasing
pressure increases
the solubility of a gas.
The relationship
that describes the
proportionality between the
solubility of a gas and the
pressure imposed by that
gas is called Henry's Law.
The proportionality
constant kH is found in
a lookup table. And that
concludes our lecture
Use your familiarity with
the common liquids to help you
make sense of solutions
and solubility.
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