Chapter 15 Notes – Solution Chemistry
Overall Summary of Chapter:
Solution chemistry has to do with what happens to compounds when they are placed in water. Ionic
compounds may dissolve and dissociate into individual ions (like NaCl becomes Na+ and Cl-) or they might
not (like PbS which stays Pbs in water). You have a table in your book on page 920 - Table C-10 which
will tell you if an ionic salt is soluble or not.
Covalent compounds (like sugar – C12H22O11) do not dissociate in water. They dissolve into individual
molecules of C12H22O11, but they do not become individual ions.
The only exception to covalent compounds are the acids (HCl, CH3COOH…etc). these covalent
compounds do separate into ions (HCl becomes H+ and Cl-. CH3COOH becomes CH3COO- and H+. They
always have a hydrogen cation (H+) and the “other” ion which is an anion (negative ions are called anions,
positive ions are called cations). Whenever an H+ is present in water, it makes a water molecule become a
hydronium ion H3O+.
Some acids dissociate better than others. HCl is a strong acid and it dissociates almost completely into
H+ and Cl- ions. But CH3COOH (acetic acid - also called, vinegar) is a weak acid and doesn't dissociate as
much into CH3COO- and H+ ions as does HCl. They both conduct electricity in water, but HCl is a much
better conductor because it separates so well.
The compounds which do dissociate into ions (soluble salts and acids), will give the water solution an
ability to carry an electric current. You could put a light bulb into a solution of salt-water, connect it to a
battery, and it would light. It would also work in an acid solution. But it would NOT work in a sugar
solution.
There are 3 big ideas in Solution chemistry that you need to know:
1. Understanding Molarity. Molarity is a fancy chemistry way of saying “concentration”. Molarity is
Moles/Liters. It is the number of moles of a substance dissolved in liters of water.
2. Knowing how temperature and pressure affect how much solute you can dissolve in water.
3. Distinguishing electrolytes from non-electrolytes. Electrolytes are those compounds which dissociate
into ions in solution. Non-electrolytes don’t. (see the overall summary of chapter above)
Vocabulary you need to know:
Substance: One element - for example: Gold, copper, silver...etc.
Solution A solid, liquid or gas which has one substance evenly distributed in another.
Solvent: The dissolving medium. (water is a typical solvent)
Solute: The substance dissolved in the solvent. (in salt-water, salt is the solute)
Homogeneous mixture: another name for a solution – usually when two solids are combined.
Alloy: A homogeneous mixture (solution) of 2 or more elements with metallic properties (ex. Brass = Cu +
Zn). Alloys are NOT chemically combined – there is no sharing of electrons.
Miscible: Like dissolves like. Non polar substances will dissolve more easily in non polar liquids. And
polar substances in polar liquids. Water is polar and will help dissolve polar compounds
Heterogeneous mixture: When you can visibly see particles in solution. (sand/water)
Solubility: The amount of a miscible substance that can be dissolved in another. See
table C-10 on p. 920 for a listing of solubilites of ionic compounds in water. This
will let you know if a single or double displacement reaction will form a precipitate or not. If it is insoluble,
that means it forms a precipitate. (Example: Sodium Chloride is soluble – NO precipitate)
Precipitate: An insoluble substance in solution.
Saturated/Unsaturated: A solution that cannot hold any more solute is saturated.
Electrolytes: Substances which dissolve into water and allow it to carry an electric current (example: Salt
water conducts electricity better than “pure” water).
Non-Electrolytes: Covalent compounds are non-electrolytes. Sugar, for example, will not dissociate into
individual ions in water – therefore it cannot carry an electric charge
A. Molarity
Concentration: Measure of how much solute is dissolved in a specific amount of solvent/solution.
 A concentrated solution contains larger amount of solute
 A dilute solution contains a smaller amount of solute
Molarity (M): moles of solute / liters of solution
 Also know as molar concentration
 It’s unit is M (read/pronounced a molar)
Example:
1) What is the molar concentration of a liter solution with 0.5 mole of solute?
Molarity = moles of solute / liter of solution
= 0.5 mol solute / 1 L solution
= 0.5 M
2) A 100.5 mL IV solution containing 5.10g glucose (C6H12O6). What is the molarity of the solution?
(MM of glucose = 180.16 g/mol)
Molarity = moles of solute / liter of solution
Step 1: convert grams of solute glucose to moles solute using MM of glucose
5.10 g glucose x 1 mole glucose = 0.0283 mol glucose
180.16 g glucose
Step 2: convert mL of solution to Liters of solution
(Remember week 1 notes: here we’re going from milliliters to liters so we’re going to a higher prefix this
means we’re moving the decimal to the left 3 decimal places)
100.5 mL = 0.1005 L
Step 3: solve for molarity
M = moles of solute/ liters of solution
= 0.0283 mol glucose / 0.1005 L solution
= 0.282 M
Homework #1: Solutions Worksheet page 1
B. How Temperature and Pressure Effect Solubility p. 460
TEMPERATURE
As the temperature of a liquid increases, the solubility of a solid in that liquid increases
EX. Hot tea will dissolve more sugar that cold tea.
As the temperature of a liquid increases, the solubility of a gas in that liquid decreases.
EX. Warm soda loses its carbonation faster than cold soda.
PRESSURE
As the pressure of a liquid increases, the solubility of both solids and gases in that liquid will
increase.
EX. When you open a can of pressurized soda, the pressure decreases and the bubbles
escape.
C. Net Ionic Equations p. 292-293
A net ionic equation lists the ionic formula for substances which are insoluble in water and just the
individual ions for those substances which are soluble in water (see p. 920 for a listing of what ions are
soluble in water and which are not). IMPORTANT – NOT ALL IONS ARE WATER SOLUBLE!
Spectator ions: Ions which are soluble in water stay surrounded by water and thus are not part of the
reaction and are called spectators. Spectator ions are dropped from the net ionic equation. See example
below where spectator ions K and NO3 are dropped.
Regular Equation: AgNO3 (aq) + KCl(aq) --> AgCl(s) + KNO3(aq)
Ionic Eq: Ag+(aq) + NO3- (aq) + K+ (aq) + Cl-(aq) --> AgCl(s) + K+(aq) + NO3-(aq)
Net Ionic Eq: Ag+(aq) + Cl-(aq) --> AgCl(s)
Homework #2: pg. 294 #33-37 AND pg. 305 #90-94 (use page 920 for help with solubility)
D. Dilutions
Dilutions: Adding additional solvent to dilute a more concentrated stock solution.
 The total # of moles of solute does not change during a dilution (only solvent is added)
Dilution Equation:
M1V1 = M2V2
M1 = the initial molarity of the concentrated solution
V1 = the initial volume of the concentrated solution
M2 = the final molarity of the dilute solution
V2 = the final volume of the dilute solution
Example:
What volume of 2.00 M CaCl2 stock solution would you use to make 0.50 L of 0.300 M CaCl2?
M1V1 = M2V2
M1 = 2.00 M
V1 = ?
M2 = 0.300 M
V2 = 0.50 L
2.00 M x V1 = 0.300M x 0.50L
V1 = 0.075 L
Homework #3: Molarity and Dilution Problems worksheet
E. Beer’s Law
Recall what we learned about light earlier in the year.
Colored solutions are absorbing some wavelengths of
light. Beer’s Law tells us that as the concentration of a
solution changes, so does the amount of light that is
absorbed.
Equation: A =  M b
A = Absorbance
 = absorptivity constant (otherwise known as the
extinction coefficient)
M = Molarity
b = path length (this is generally 1)
E. Molality
m = molality = # moles of solute/kg of solvent
Molality is used to predict temperature changes between boiling point and freezing points of liquids.
DON’T GET THIS CONFUSED WITH OUR FRIEND, MOLARITY, M with is Moles/liter of H2O
Practice Problem: How would you prepare a 0.50 m solution of NaCl in 500 g of water?
0.50 moles
1 kg
=
x moles
0.5 kg
x = 0.25 moles of NaCl
convert moles to grams and add to 500 ml of water
* remember water’s density is 1 g/ml so 1000 ml is 1000 grams and 1000 grams = 1 kg
E. Colligative Properties pg. 471
Colligative properties - A property of solutions that doesn’t depend on the size or type of molecule or
atom in solution, just the concentration. For example, the Freezing point and boiling point of a liquid are
determined by the number of particles in solution, not the type. 1 mole of sodium particles in 1L of water
will do the same as 1 mole of sugar in 1L of water. This is very important.
F. Calculating Freezing Point Depression and Boiling Point Elevation pg 472-475
Tb = kb x m x i
Tf = kf x m x i
T is the change of temperature from what it would normally be without any solute. When dealing with
boiling, the Tb would be added to the original boiling point. When dealing with freezing, the Tf would be
subtracted from the original freezing point.
m is molality
K is a constant used for the solvent you are talking about changing the boiling of freezing of.
For example: kb for water is: 0.51 oC(kgH20))/mol
kf for water is: 1.86 oC(kgH20))/mol
i is the number of ions that the solute dissolves into. For example, NaCl dissolves into two ions – Na+ and
Cl- so i = 2 for NaCl. Aluminum nitrate (Al(NO3)3) dissolves into one Al+3 and three NO3- so i = 4.
See Table 15-4 & 15-5 on pg. 472 & 474 for more k values of different solvents See Table C-12 on pg. 921
Fact: By adding a solute to any liquid, you increase the boiling point and reduce the freezing point.
For example: adding salt to water will increase its boiling point and decrease its freezing point.
Why does it increase the boiling point? You lower the vapor pressure of the liquid when you add a solute
because the solute prevents the pure solvent molecules from escaping. (Recall that vapor pressure is the
internal desire of a liquid to boil away). A lower vapor pressure means you have to heat water hotter to get
it to vaporize. Salt water will boil at a temperature greater than 100oC.
Why does it lower the freezing point? By adding a solute to a liquid solvent, you decrease the amount of
intermolecular interactions between molecules of the pure solvent. Thus, if pure water can be frozen at 0oC,
salt water will freeze at a colder temperature because the salt ions get in the way of the intermolecular
attraction between water molecules. When salt water freezes, the frozen substance contains water only, not
salt water – want proof? Icebergs floating in the ocean are pure water!
More proof? How about this one: Spreading salt on snow melts it because the outside air temperature is
cold enough to freeze pure water, but not cold enough to freeze salt water.
What does this have to do with making ice cream? Ice cream is mostly frozen milk and sugar. You can
think of ice cream as mostly water with some dissolved solute in there (lactose, protein…etc). This means
that ice cream will freeze at a temperature BELOW 0oC (the normal freezing point of pure water.). Your
freezer is kept at about –10oC. Cold enough to keep ice cream from melting.
Pure ice (made from water) will draw heat out of the surrounding air to begin melting into liquid water. As
ice melts, its temperature remains constant until all of the ice has melted (we will learn more about this
interesting phenomenon of solids keeping their temperature constant until they have completely melted in
Chapter 16 – Thermodynamics) Normally, it will draw in enough heat to reach its melting/freezing point
of 0oC. BUT…..If you add salt to the ice, this reduces the melting/freezing point of ice to about –10oC
(remember adding a solute to a solvent will reduce the freezing point). Now, the ice will draw in enough
heat to make its temperature –10oC. This is cold enough to freeze milk into ice cream.
WOW, This is cool: Did you know that adding 1 mole of CaCl2 to water will drop the freezing point
lower than if you added 1 mole of NaCl? Why is that? Because 1 mole of CaCl2 will break up into 1 mole
of Ca+ cations and TWO moles of Cl- anions. But NaCl will break up into only 1 mole of Na+ ions and 1
mole of Cl- ions. The colligative property rule states that more ions makes a bigger change in the freezing
point and boiling point. 1 mole of CaCl2 essentially becomes 3 moles when it breaks up. 1 mole of NaCl
becomes 2 moles.
How many moles does 1 mole of C12H22O11 become when dropped into water? 1 mole. Why?
Because it doesn’t break apart!!!!!!
Homework #4: Practice Problems # 33-36 p. 474
Pg 485 #76-79, 83, 86-88