Chapter 13 Notes – Solution Chemistry
1. Definitions
of Solutions:
Substance: One element - for example: Gold, copper, silver...etc.
Solution: A solid, liquid or gas which has one
substance evenly distributed in another.
Miscible (Immiscible): Like
dissolves like. Non polar substances
will dissolve more easily in non polar liquids. And polar substances in polar liquids (p. 406)
Hydration: The process
of polar water molecules pulling ionic compounds apart so that their ions float
around When you crystallize compounds which were once in water, you get
“hydrates”. Hydrates have a specific
ratio between the number of water molecules and ionic compounds. (p. 405).
Homogeneous mixture: another name for a solution –
usually when two solids are combined.
Alloy: A homogeneous mixture of 2 or more elements with
metallic properties (ex. Brass = Cu + Zn)
Heterogeneous mixture: When you can visibly see particles in solution. (sand/water)
Solvent: The dissolving medium.
Solute: The substance dissolved in the solvent.
Solubility: The amount of a miscible substance that can
be dissolved in another. See
table A-12 on p. 900 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.
Suspensions: If the
particles in a solvent are so large that they settle out unless the mixture is
constantly stirred (example: Muddy
water)
Colloids: Particles
that are intermediate in size between those in solutions and suspensions. (Example:
The “cloudy” part of muddy water left over after the big chunks settle
to the bottom. Shaving cream, mayonnaise,
smoke, fog, cheese, butter)
Electrolytes: Substances
which dissolve into water and allow it to carry an electric current
(example: Salt water conducts
electricity better than “pure” water).
See example on p. 400
Solution Equilibrium: Where
dissolution and crystallization are occurring at the same rate (p. 402)
Saturated solution: Contains the
maximum amount of solute in a solvent. (p. 403)
Supersaturated solution: Contains more
than the maximum. This is unstable and
if “seeded” crystals will form (p. 403)
Read the section: Heats of Solution (p. 409-410) aloud in class
2. Changing
the Rate of Dissolution:
1. Increasing the Surface area
of the solute (p. 401)
2. Agitating the solution (p.
401)
3. Heating the solvent (p. 402)
4. How Temperature and Pressure
Affect Solubility? (p. 407-409)
Temperature: As the temperature of a liquid increases,
the solubility of a solid in that
liquid increases. Example. Hot water will dissolve more sugar than cold
water.
Temperature II:
As the temperature of a
liquid increases, the solubility of a gas
in that liquid decreases. Example:
Warm soda will lose its fizz but a
cold soda won’t (as fast).
Pressure: As the pressure on a liquid increases,
the solubility of both solids or gases in that liquid will increase. (pressure doesn’t affect it very much, but it does increase
solubility a little bit) (p. 407)
Example: When you open a can of soda to the reduced
pressure of outside air, the bubbles escape.
1. Net Ionic
Equations – (p. 429)
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. 900
for a listing of what ions are soluble in water and which are not). IMPORTANT – NOT ALL IONS ARE WATER
SOLUBLE! There are some general rules for solubility
on p. 427 Table 14-1
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 usually
dropped from the net ionic equation. See example below where spectator ions K
and NO3 are dropped.
Ionic Eq: Ag+(aq)
+ NO3- (aq) + K+ (aq) + Cl-(aq) --à AgCl(s) + K+(aq) + NO3-(aq)
Net Ionic Eq: Ag+(aq) + Cl-(aq)
-----à AgCl(s)
Making Crystals: Crystals are made from changes
in solubility due to temperature. You
can dissolve a lot of sugar in hot water, but not in cold. So, you supersaturate a hot water solution
with sugar and then slowly cool it to bring out the crystals.
2. Ionization
(p. 431-433)
When you put ionic compounds
into water they will do one of three things:
1) nothing 2) dissolve somewhat
3) dissolve a lot! Electrolytes are
dissolved ionic compounds in water which can then carry an electric charge
(remember chapter 13). Weak
electrolytes don’t dissolve much (such as HF).
But strong electrolytes (like NaCl mostly break apart completely in
water). COVALENT MOLECULES USUALLY NEVER BREAK APART LIKE IONIC ONES DO!!!!!!!! (an
exception is the hydrogen halides like HCl, HBr and HI which are covalent and
do break apart)
Question: AgCl is practically insoluble in water and yet it is considered
to be a strong electrolyte. What’s the
deal? The deal is that the small part
of AgCl which actually does dissolve in water is completely dissociated into
Ag+ and Cl- ions which can then carry an electric charge. (p. 432-433)
3. Freezing
point/Boiling Point Changes - AKA Colligative Properties. (p. 436-441)
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 melting
point of a liquid are determined by the number of particles in solution, not
the type. 1 mole of salt in 1L of water
will do the same as 1 mole of sugar in 1L of water.
Ions vs. molecules in determining F.P and B.P.
changes:
Ions separate into more than one particle in water. For example, NaCl breaks up into Na+
and Cl- ions in water. But C12H22O11
(sucrose) remains in its molecular form in water. Therefore, 1 mole of NaCl makes 2 moles of solute in water and 1
mole of sucrose makes 1 mole of solute in water. Calcium Chloride (CaCl2) turns into 3 ions and is
therefore triple the effectiveness as sucrose in changing b.p and f.p.
Ions in solution are called electrolytes because they can make a solution carry an electric
charge due to the charges on the ions. (see p. 431-433)
5. Molality:
(p. 416-418) m = molality = # of moles of solute/Kg of
solvent.
Molality is used to predict temperature changes
between b.p. and f.p. of liquids.
REMEMBER WHEN…….
M = Molarity = moles of solute/L of solution (which would
include L of the solute if applicable).
Molarity is used do determine the number of moles of a substance in
solution.
Molality problem: How would you prepare a 0.50 m
solution of sucrose in 500.0 g of water?
Remember that the density of
water at room temp is about 1g/mL or 1 g/cm3. Therefore, if you have 1000 mL of water, you
also have 1000g or 1Kg.
6. 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.
7. Calculating
F.P. depression or B.P. elevation:
DTb = kb
x m x # of particles (see VERY IMPORTANT below to understand “# of
particles”)
DTf = kf
x m x # of particles
k is a constant used for the solvent you are talking about –
for example, water. (see table 14-2 p.
438)
Kb for water is = 0.51 oC(kg H2O)/molsolute
Kf
for water is = -1.86 oC(kg H2O)/molsolute
VERY IMPORTANT:
Molality of a solution needs to be doubled or tripled when you are using
an ionic substance as your solute. For
example, a molality value of 1m for a solution of NaCl is actually 2m. But a molality value of 1m for a solution of
sucrose is still 1m. NONELECTROLYTE
solutes (such as insoluble ionic compounds or insoluble covalent compounds) do
NOT break apart!
8. Alloys:
An alloy is a homogeneous mixture of two or more elements to
make a substance that is metallic in nature.
For example, if you mix copper and zinc together you will get
brass. But there is no chemical bond
created between the zinc and the copper and therefore brass is NOT a chemical
compound (like water or salt or sugar…etc).
It is an alloy.
Alloys generally have lower electrical conductivity than their
original metals. They also generally
have a lower melting point. Generally,
alloys are stronger than their original metal and are less likely to
corrode. Pound for pound, alloys are
stronger than their original metals and therefore a smaller amount of alloy can
be used than an original metal.
There are two basic types of alloys: Substitutional alloys and Interstitial alloys:
Substitutional alloys are where one metal atom is
replaced by another (example: brass)
Interstitial alloys are where one metal has another
atom in the spaces between its atoms
(example:
steel)
Example uses of alloys:
Steel is an alloy of carbon
and iron. Uses: Making cars. Steel is stronger than either carbon or iron. So a thinner amount
of steel can be used in the door of your car than a piece of iron. Also, steel will not rust as easily as pure
iron.
Special metal alloys are used
in spacecraft because they are lighter, stronger and more heat resistant than
their original elemental parts.
Brass is used in faucet
fixtures because it will not corrode as easily as a zinc or copper fixture
(remember brass is an alloy of zinc and copper).
Bronze is used in making
statues because it is more bendable and corrodes less than its original atoms
(which are Cu, Zn, Pb and Sn).
Solder is used to hold
electrical contacts together because it has a lower melting point than its original
atoms (which are Pb and Sn). You can
use a soldering iron to melt solder, but it won’t melt Pb or Sn.
Sterling silver (and likewise
stainless steel) is used to make silverware which won’t tarnish as much as pure
silver. Sterling silver is made out of
Ag and Cu.
Dentists use a special alloy
called an “Amalgam” which is an alloy of Hg and Au for fillings. Pure gold is not as strong or as pliable as
it is when mixed with Hg. Any alloy with
Hg is called an amalgam.
Homework for Chapter 13 and 14:
20 points Total
Predicting Solubility of
compounds: p. 447 #2
Net Ionic Equations: Practice
Problems 1-4 p. 430; p. 448 #13, 16, 17, 18
Molality: Practice
Problems # 1-4 p. 418, p. 421 # 24, 25
Molarity (our old friend): Practice Problems p. 415 #1-3; p. 421 # 17, 18 (skip 18b), 20
Freezing Point/Boiling
Point Change Problems: Practice Problems 1-4 p. 440; Practice
Problems 1-4 p. 441; Practice Problems 1-3 p. 445; p. 449 #19, 20, 25, 26, 27,
28, 30
VERY IMPORTANT: Molality of a
solution needs to be doubled or tripled when you are using an ionic substance
as your solute. For example, a molality
value of 1m for a solution of NaCl is actually 2m. But a molality value of 1m for a solution of sucrose is still
1m. NONELECTROLYTE solutes (such as
insoluble ionic compounds or insoluble covalent compounds) do NOT break apart!