Solutions
Chapter 15
Mr. O’Brien (SFHS)
Solutes & Solvents (std.6a)



What are solutions?

A homogeneous mixture of two or more
substances.

homogeneous: has a uniform composition
(looks the same throughout).

Can be any state of matter.

GAS: air (N2+O2+Ar+CO2)

LIQUID: chocolate milk (Milk & Chocolate)

SOLID: steel (Fe & C)
What are solutions composed of?

Solute: The “stuff” that gets dissolved.

smaller amount

Solvent: Dissolves the solute.

larger amount

water is a universal solvent
What are not solutions?

Pure elements/compounds

Heterogeneous mixtures

mixtures that do not form a uniform
composition (two or more layers)
 Examples: oil & water, Italian dressing
(figure 2) Above shows examples of a
heterogeneous mixtures. Why do these
mixtures have two or more layers?
(figure 1) Above shows examples
of a solution. In chocolate milk,
water is your solvent and chocolate
mix the solute. In air, Nitrogen is the
solvent and Oxygen is the solute.
The Dissolving Process (std.6b)

How do particles dissolve in each other?


The RANDOM MOTION of solute particles colliding
with solvent particles explains the dissolving process.
Why can water dissolve salt but not oil?

“Like dissolve like rule”


Compounds that contain charges (POLAR) can
dissolve in each other –OR- compounds that contain
non charges (NON-POLAR) can dissolve each other.
Compounds that are opposite cannot dissolve
each other.


(figure 1) The green dye will slowly dissolve
in the water because of its random motion. If
you heat up the solution then the dissolving
will occur faster because increased kinetic
energy.
Oil (nonpolar) & water (polar) = no solution
NaCl (ionic) & water (polar) = solution

remember ionic compounds contain charges.
Solvent-Solute Combinations
Solvent Type
Solute Type
Is Solution Likely?
Polar
Polar
Yes
Polar
Nonpolar
No
Nonpolar
Polar
No
Nonpolar
Nonpolar
Yes
(figure 2) The image above shows two nonpolar
substances dissolving (green solution). The right
image shows a heterogeneous mixture. Because
the mixture is composed of polar water and
nonpolar hydrocarbon (oil), it forms two layers.
The Dissolving Process (std.6b)

How do attractive forces play apart?





If solute is able to dissolve:
then the force that holds the solute together is
WEAKER than the force between the solvent
and solute.
If the solute does not dissolve:
then the force that holds the solute together is
STRONGER than the force between the solute
and solvent.
How does salt dissolve in water?
1.
2.
3.
4.
Water (solvent) COLLIDES with salt (solute)
STRONG ATTRACTIVE FORCES between
solvent and solute then separate the salt.
(H+) ions of water are ATTRACTED to the (Cl-)
ions.
(O-) ion of water are ATTRACTED to the (Na+)
ions.
(figure 3) The figure to the left shows the dissolving
process of NaCl (salt) in water. Salt is able to dissolve
because the attraction of water molecules to the salt
ions is much stronger than the attraction between the
salt (NaCl) itself. Note the orientation of water around
Na+ and around Cl-.
There is a limit of how much solute
can dissolve in a given solvent.
(figure 1) In beaker (a) and (b) the solute is being
dissolved. In (c) the solution has reached maximum
saturation because it’s the same as (d)
EXTRA info



Unsaturated solution: solvent is still able
to dissolve solute.
Saturated solution: solvent is unable to
dissolve any more solute.
Supersaturated solution: contains more
dissolved solute than a saturated solution
@ same temp
Factors that Affect Solubility (std.6c)

What is solubility?


Maximum amount of solute that can be dissolved
in solvent.
What are three factors of solubility?
 Surface Area

By breaking up larger solid pieces in to smaller ones.
More surface area → more collisions between solute &
solvent = increase solubility.

Temperature

In solids (direct relationship) ↑ temp. ↑ solubility.

In gases (indirect relationship) ↑ temp. ↓ solubility.

Pressure

The solubility of gases in liquids increases as pressure
increases.

(figure 1) Notice that there are two types of
lines. In the gas state, as you increase temp. the
solubility decreases. But in solids, it is the
opposite.
(figure 3) LEFT. Gas (CO2) is dissolved in soda. A warm soda will lose its “fizz”
faster than cold soda because gases dissolve in liquids better in colder temperatures.
Gases will slow down in colder environments. RIGHT. The more pressure you apply
to a gas the more it will dissolve in a liquid. Similarly when you open a soda can, you
are releasing the pressure so the gas will naturally want to escape.
(figure 2) Explanation of increasing surface area
Factors That Affect Solubility (std.6c)
Using the two graphs
answer the following:
Explain why the solubilty graphs of HCl, NH3, and
SO2 are opposite the rest of the compounds.
1.

HCl, NH3, and SO2 are gases. As temperature
increases they become less soluble. The others are
solids.
Which compound has the highest solubility at 0oC?
Which compound has the lowest solubility at 0oC?
2.


Highest solubility @ 0oC = KI
Lowest solubility @ 0oC = KClO3
What could be said about the solubility of NaCl
compared to the other solvents as the temperature
increases?
If the temperature increases to 110oC which solute
will have the highest solubility?
3.
4.

KNO3 has a larger slope than any other salt.
At 100oC 40g of NaCl can dissolve in 100g of H2O.
How much can dissolve at the same temperature
with only 50g of H2O?
5.

20g of NaCl can be dissolved because there is half the
amount of solvent available (50g).
Calculating
Concentration (std.6d)

Concentration: a measure of
how much solute is dissolved in a
specific amount of solvent or
solution.



A concentrated solution contains
a lot of dissolved solute.
A dilute solution contains little
dissolved solute.
(figure 1) Shows concentration of sugar in various drinks.
The following are ways
concentration expressed:
solute X 100
Percent by mass
=
or volume
solution
parts per million (ppm) =
density =
mass of object
volume of object
solute
X 106
solution
molarity =
(figure 2) Dilute solutions compared to concentrated
solutions. The dots represent dissolved solute.
moles of solute
liters of solution
Calculating Concentration
(std.6d)

Using Percent & PPM to Describe
Concentration

In Ortho insecticide, there is 25g of the active
ingredient Diazinon in 75g of water. What is the
percent by mass of Diazinon in the solution?
25g solute
X 100
100g solution

= 25%
A pond was found to contain 50g of mercury.
The estimated mass of the pond is 200kg
(200,000g). What is the concentration in ppm?
50g mercury
200 000 ponds mass solution
X 106
=250ppm
(figure 1) LEFT. An actual picture of
Ortho insecticide showing concentration
expressed in %. RIGHT. Notice how
small ppm concentration are.
(figure 2) Crest toothpaste mentions a concentration of
sodium fluoride (0.243%) which translate to 0.243g of
NaF per 100mL of toothpaste. It also mentions a
concentration (0.15%) which translate to 0.15 grams of
fluoride ion per 100mL of toothpaste.
Calculating Concentration (std.6d)

Practice:

Blood contains .01g of calcium ions in 100g
of blood serum. What is this concentration
in ppm?
+2
6
.01g Ca
X 10
100g blood

Calculate the percent by mass of 15g of
NaCl dissolved in 45g of H2O.
15g NaCl
X 100
45g solution



If I add 200g of water to 50.0g of sodium
acetate, what is the percent by mass of
sodium acetate in this solution?
Calculate the mass of a solution if the mass
of a solute is 25g and the percent of solute in
the solution is 25%.
[HONORS] What mass of water must be
added to 200g NaCl to make a 20% by mass
aqueous solution?
200g solute
800g solution = 200g + ?gH2O
?g solution
50g sodium acetate
250g solution
25g solute
?g solution
X 100
= 20%
X 100
X 100
Calculating Concentration
(std.6d)
 Molarity (M) Most common concentration
unit.
uses moles (moles allows us to count molecules)


Example:
A 100mL intravenous (IV) solution contains 90g
of glucose (C6H12O6). What is the molarity of this
solution?

1. Convert (g) to moles.
90g C6H12O6
1mol C6H12O6
1
180g C6H12O6
2. Convert (mL) to L.
100mL
1L
= 0.1L
1
1000mL
= 0.5mol C6H12O6
3. Substitute and solve
0.5mol C6H12O6
(figure 1) Each of these solutions contain a different
amount of solute. The amount of molecules or mass can
be calculated because molarity is based on moles.
= 5M
0.1L solution

Practice:

How many moles are there in 1L of a 0.35M of
moles
= .35M
MgCl2?
1L solution

A liter of 2M NaOH contains how many grams of
moles
= 2M
NaOH?
1L solution

What is the molarity
(M) of
9g H2O in 1000mL?
0.5 moles
=M
1L solution
(figure 2) Look here for a quick review on mole
conversions. In molarity problems you will be mostly
converting between moles and grams. Use molar mass.
Colligative Properties (Honors)

Colligative. “depending on the
collection.”

Physical properties of solutions that are
affected by the number of particles NOT
type of solute.


Boiling point:

Boiling pure solvent requires less heat
than boiling a solvent with a solute
(solution)



Such as boiling point and freezing point.
Ex: pure water boils at 100oC but with
NaCl water can boil at 200oC.
More the solute particles added =
INCREASE in boiling point.
Freezing point:

Freezing point of a solution is
ALWAYS LOWER than that of a pure
solvent.

Ex: pure water freezes at 0oC BUT with
NaCl water can freeze at -50oC.