UNIT 6 - Solution Chemistry
Lesson 1 Solutions and Solubility – This is a “flipped” intro – so what that means is
that you watch the video and take notes, for homework. https://www.youtube.com/watch?v=f5a9xPDreRY
Watch: Solutions
Characteristics of solutions:

solute - the material that is being dissolved into
something else.
Example : when we dissolve sugar into a cup of coffee,
the sugar is the solute.

solvent - the material into which a solute is being
dissolved. In the above example, the coffee is a
solvent. Solvents can be almost any material or in any
state but for the most part, we will be using water as a
solvent.

solution - a combination of a solute and a solvent. We
can not distinguish the identify a solute and a solvent- it is a HOMOGENEOUS MIXTURE!

A solution can be in any combination of states:
- a solid/liquid combination (dissolving a solid in water such as salt dissolved in water).
- a liquid/liquid combination (table vinegar is a mixture of water and acetic acid).
- solid/solid combination (amalgam or any type of alloy).
- Air is an example of gas solution: nitrogen
78.08%
oxygen
20.955
argon
0.93%
Carbon dioxide 0.03%
-solutions having water as the solvent is referred to as AQUEOUS solution (blood, saliva, etc).

There comes a point where no more solute can dissolve into the solution. At this point in time the solution is said to be
saturated.
- When solute can still dissolve in a solution, it is said to be unsaturated.
- To increase the solubility of the chemical, we can usually heat the solution up.
- The result will be a solution that contains more solute than what would normally be dissolved in the
solvent. This is called a supersaturated solution.

Solubility of a solute is the maximum amount of solute that can dissolve in a given amount of solution at a given
temperature.
-
Solubility is temperature dependent. Usually an increase in temperature increases the solubility and
increases how fast a solid dissolves in water.
Solubility is measured in terms of g/L, g/mL, or mol/L (at a given temp.)
Solubility of solids in liquids vary widely:
Does the solubility of sodium nitrate increase or
decrease with an increase in temperature?
Which compound has a solubility of 61g / 100g H2O at
70C
Which compound has a solubility of 41g/100g H2O at
50C
How can we increase solubility?
Check for Understanding:
Solubility solids /
liquids
pressure
increases
Pressure
decreases
Temperature
increases
Temperature
decreases
-
Solubility of
gases
1.
In soda, CO2 is the ____ and water is the ____.
2.
How would lowering the pressure change the
fizziness of a soda?
3.
Which term describes a solution in which only a small
amount of solute is dissolved?
There is an upper limit to the solubility of a solid. It important to note that even the least soluble solid will
have a few particles dissolve per litre of solution.
So, we will say that a substance is soluble if 0.1 mol of the substance dissolves per litre.
When there is no apparent limit to the solubility of one substance in another, the components are said to be MISCIBLE.
THIS IS THE END OF YOUR FLIPPED LESSON. SEE YOU IN CLASS TOMORROW
Let’s Make some Rock Candy – Super-Super-saturation https://www.youtube.com/watch?v=HvKJz4M585c
Assignment: Write out the procedure that you followed in order to make your Rock Candy. For each step, explain what
is happening and why. It should be put into a table like the one below.
Procedure
Explanation
Electrical Conductivity:






Metals (solid and liquid phase):
Non-metals (solid and liquid phase):
Ionic compound in solid phase:
e.g. NaCl(s),
Ionic compound in liquid and aqueous phase:
Covalent compounds in solid and liquid phase:
Acid and Base:
conduct electricity
non-conductor of electricity
non-conductor
conductor
non-conductor (usually!)
conductor
Electrolytes and Nonelectrolytes





An electrolyte is a substance that will conduct electric
current when it is dissolved. It will conduct electric current
because it exists as charged particles (ions) when it is
dissolved. Good examples: ionic compounds (such as
NaCl), acids, and bases.
A nonelectrolyte is a substance that will not conduct
electric current when it is dissolved. It will not conduct
electric current because it exists as uncharged molecules
when it is dissolved. Good example: table sugar (sucrose,
C12H22O11).
Note: Ca2+, Na+, and K+ are the primary electrolytes in
your body that are responsible for conducting electricity.
Your body uses a lot of electricity to function – for
instance, every time you move a muscle! The electricity
does not travel through wires, though. It travels through your nerves.
Water, a molecular compound, has very few charged particles in it when it is pure. It has a natural concentration
of H3O+ and OH- ions of 0.0000001 M! Other than that, it doesn’t contain any ions. Thus, pure water is a poor
conductor of electricity.
The electrolyte will allow the bulb to light up when it’s dissolved in solution. The nonelectrolyte will not.
Examples: Which of the following would you expect to form conducting (ionic) solutions?
a) KI
b) ICl
c) HBr
d) CH4
e) LiOH
f) CH3OCH3
g) N2O
h) HCl
Examples: Which of the following would you expect to conduct electricity?
a) NaCl(s)
b) CH3OH(l)
h) KBr(aq)
i) KBr(s)
c) Cu(s)
j) Ag(s)
d) Fe(l) e) K2CrO4(aq)
k) Hg(l)
f) C14H10(s)
g) CO2(s)
l) Na(s)
Hebden Page 198 #6-8
Crash Course Chemistry: https://www.youtube.com/watch?v=AN4KifV12DA&index=7&list=PL8dPuuaLjXtPHzzYuWy6fYEaX9mQQ8oGr
Molarity
Vs
END OF LESSON
Molality
Polarity of Molecules:
The attractive forces that exist between molecules are known as intermolecular forces. These include ionic interactions, dipole-dipole
interactions and dispersion or London dispersion forces. Dipole-dipole interactions and dispersion forces are weaker than thermal
energy (2.4 kJ/mole) at room temperature and are referred to as Van der Waals Force.
A. Van der Waal’s Forces (intermolecular interaction):
1.
Dipole-dipole interactions:
2. London Forces:




molecules that are dipolar are attracted
electrostatically with one another.
Polar molecules have a partial positive and a partial
negative portion of the molecule.
They are made of atoms having differences in
ELECTRONEGATIVITY.
Polar molecule must also be ASYMETRICAL.
https://www.youtube.com/watch?v=3t1Jn_jrsQk

are weak attractive forces which arise as the result
of temporary dipolar attraction between
neighbouring atoms.
 London Forces are always present, even in species
that have dipolar attraction.
 This is a weaker attraction than dipole-dipole forces.
B. Hydrogen Bonding (intermolecular interaction):




Hydrogen bonding occurs when hydrogen is bonded covalently
to N, O, or F (high electronegativity). Highly polar.
Hydrogen bonding is stronger than the Van der Waal’s Forces.
The presence of H-O- , H-N-, H-F indicate a highly polar
section of the molecule.
Asymetrical molecules are POLAR molecules.
*****Hydrogen bonds are not actually bonds at all, they are intermolecular attractions !!!
Hebden Page 198 #9-16
Electronegativity of elements


Covalent bonding is the sharing of electrons.
Unless the two atoms sharing the molecules are identical, the sharing will not be equal

The more electronegative atom will attract the electrons more strongly than the less electronegative atom.
 This will result in the more electronegative atom having a slight negative charge (-)
 The less electronegative atoms will be slightly deficient in electrons and thus have a slight positive charge (+)
This type of covalent bond in which atoms have slight electrical charges are called polar bonds
Eg.


Electronegativity of elements can be judged from their position on
the periodic table.
All elements involved in covalent bonds do have quite high
electronegativities.
B & Si < P & H < C & S & I < Br < Cl & N < O < F
High

Very High
Extremely High
The greater the difference in electronegativity of the atoms involved
the greater the polarity of the bond.
In many molecules, the polar bonds result in the molecule have a resultant dipole
i.e. there is a positive and negative end to the molecule
 In some molecules, the effects of the polar bonds cancel out because of the symmetry of the molecule.
For a molecule to be polar


It must contain polar bonds
Its shape must be such that the centre of positive and negative charges are not in the same place
Shapes of molecules Take Notes while watching “ VESPR: Theory Introduction” Video by
Tylar DeWitt on You Tube. Video is on the wvssearland website. (https://www.youtube.com/watch?v=nxebQZUVvTg)
VSEPR:
What is “inside” s covalent bond?________________________________
No. of
“things”
around
central
atom
Bond
Paris
and
Lone
Pairs
Bond
Angles
2
2 bp
180
LINEAR
3
3bp
120
TRIGONAL PLANAR
2bp +
1lp
116
BENT
4bp
109.5
TETRAHEDRAL
3bp +
1lp
107
TRIGONAL PYRAMIDAL
2bp +
2lp
105
BENT
4
Why is there Repulsion?_______________________________________
https://www.youtube.com/watch?v=keHS-CASZfc



The shapes of molecules are determined by the repulsion between the electron pairs in the valence level.
This is known as the Valence Shell Electron Pair Repulsion Theory (VSEPR Theory)
In the most common molecules, filled valence levels contains four pair of electrons.
Some molecules also contain non-bonding or lone pairs of electrons. These lone pairs affect the shape of the molecule.





In ammonia, NH3, there are three pairs of bonded electrons and one lone pair.
Without the lone pair, the molecule would look like an equilateral triangle with a nitrogen atom at its centre.
The lone pair that is not involved in bonding repels the bonding electrons so the molecule has the shape of a
trigonal pyramid.
In water, there are two pairs of bonded electrons and two lone pairs.
The two lone pairs cause the molecule to be bent, not linear as would be expected without the lone pairs.
Resonance Structures
Often two or more equivalent structures can be drawn for a molecule. These are known as resonance structures.
Resonance refers to the arrangement of valence electrons in molecules or ions for which more than one Lewis structure can be
written. The actual molecule that exists is often referred to as a resonance hybrid of these structures.
Eg.
O3
Resonance does NOT mean than the molecule flips from one structure to another. The bonds actually have lengths and
strengths intermediate between those of single and double bonds or double and triple bonds.
Steps for writing Lewis structures
1. Write the correct arrangement of the atoms using single bonds. Where necessary, apply the following guidelines.
a) Smaller, more electronegative non-metal atoms surround larger, less electronegative non-metal atoms.
b) Oxygen, hydrogen, and/or halogen atoms often surround a central metal or non-metal atom in a symmetrical
arrangement.
c) Carbon atoms are usually bonded to each other.
d) Oxygen atoms are bonded to each other only in peroxides (or superoxides)
e) In most acids, such as H2SO4, and in many other compounds that contain both oxygen and hydrogen atoms, the hydrogen
atoms are all bonded to oxygen atoms.
2. Find the total number of valence electrons. Add together the number of valence electrons contributed by each atom.
If the species is an ion, subtract one electron for each unit of positive charge or add one electron for each unit of negative
charge.
3. Assign two electrons to each covalent bond.
4. Distribute the remaining electrons so that each atom has the appropriate number on nonbonded electrons. For
elements from the second period, other than beryllium and boron, this is the number of electrons needed so that each
atom is surrounded by an octet. For elements of the third period and beyond, except aluminum, this is often the number
of electrons needed to complete on octet, although extra electrons can also be placed around atoms if these elements
when they are the central atoms in compounds. Remember that atoms bonded to a central atom usually obey the octet
rule.
5. If there are not enough electrons to go around, change some of the single bonds to multiple bonds. Multiple
bonds can be written between carbon, nitrogen, oxygen, sulphur, selenium and phosphorous atoms. (Note that beryllium,
boron, and aluminum do not form multiple bonds.
Steps in using VSEPR to predict geometry
1. Write the Lewis structure of the molecule.
2. Determine the number of bonding pairs and only pairs of electrons around the central atom.
3. Determine the ideal geometry. Then, if necessary taking into account the presence of lone pairs, predict the actual shape
of the molecule
4. Keep in mind that lone pairs occupy large site (equatorial in molecules derived from AB5 molecules) or, when sites are
equal, occupy sites opposite rather than next to each other.
VSEPR Worksheet
END OF LESSON 2
Polar and Non-Polar Solvents:
Polar Molecules:
Solvent
- Dipole is present
- Molecule is ASYMETRICAL (not symmetrical).
Hebden p. 204-205 – complete the table below
Molecular
drawing
Polar or
NonPolar
Water
Solvent
Methanol
Acetic acid
Ethanol
Benzene
Chloroform
Ethoxyethane
Carbon
tetrachloride
Acetone
Heptane
Liquid
ammonia
Molecular
drawing
Polar or
NonPolar
“Like dissolves like” :

polar solvents dissolve polar solutes and ionic compounds
 Non-polar solvents dissolve non-polar solvents
EXPLANATION:
The “dissolving process” involves three different attractions
 Solvent molecule  other solvent molecules
 Solvent molecule  solute particles
 Solute particles  other solute particles
Solvent - Solvent
Solvent - Solute
Solute - Solute
Energy must be added to break a bond and is released when bonds form. If Solvent - Solute bond
energy is > Solute - Solute, then the nrg from forming the new bond can be used to break the old
one. Strong Solvent – Solvent then allows solutes to be mixed within the solvent = dissolving
Explanation
Conclusion
Solubility of Polar & Ionic
Solutes
Solubility of Nonpolar Solutes
A. Polar and ionic molecules are soluble in polar
solvents.
ex:
Ex:
Salt (ionic) dissolves in water (polar solvent).
I2(s) is more soluble in water than gasoline.
non-polar molecules are soluble in non-polar solvents.
B. Polar and ionic compounds are generally held
together quite tightly.
Polar solvents are strongly attracted to the ions or the polar
compounds causing separation of the solute out of its crystal
structure and into solution.

Non-polar solvents do not have this attraction, therefore, polar
solutes are unable to separate into a non-polar solution.
C. Non-polar solutes are held together by London
Forces.
Solvents with enough London Forces to attract the solutes
and break off the solid crystal structure are capable of
dissolving the solid. Generally speaking, non-polar solvents
have high London Forces, and dissolve other non-polar
compounds.
Hydrogen bonding holds water molecules together. For example, water molecules undergo
hydrogen bonding. When water freezes, the six sided structure of water crystals is due to
hydrogen bonding.
Hebden p. 207 # 18-22
Read – 207-208 then do #23-27
END OF LESSON 3 – QUIZ #1 UP TO THIS POINT!
The nature of Solutions of ions:

The formation of solution depends on the ability of the solute to dissolve in the solvent.



Ionic solid are crystal structure made up of ions.
Covalent solids are crystal structure made up of covalent molecules.
Crystal lattice is an orderly arrangement of particles that exists within a crystal.

Water: “Universal Solvent”, polar, V-shaped molecule. Its positive end attracts the negative ions and the negative
end attracts the positive ions.

Dissociation: reaction that involves the separation of ions in an ionic compound.

Ionization: reaction that involves the breaking up of a neutral molecule into ions.
Read p. 133-148, Zumdahl.
Do Exercises: p. 180-181 # 11, 13, 15, 17, 25
Hebden p210 #28-29
END OF LESSON 4
Calculating The concentration of Ions in Solution:
concentration - this is the amount of solute that is in a certain amount of solution. Because 'amount
of solute' can be interpreted in a number of different ways (i.e. grams, moles, molecules etc.) we will
define a value for concentration :
This is also defined as MOLARITY.
The units for molarity are moles/litre, mol/L or M for short.
Examples :
1) What is the concentration of a solution containing 0.055 moles of HCl dissolved in water to make
250.0 mL of solution
2) What is the concentration of a solution containing 0.055 grams of HCl dissolved in water to
make 250.0 mL of solution
3) How many molecules of NaOH are dissolved in 125.3 mL of a 3.00 M NaOH solution ?
How to make a solution of a certain concentration:
In order to make a solution of certain concentration, we do NOT dissolve the amount of
solute needed in the volume of solvent. If we did that, the volume would increase somewhat and we
would not have the precise concentration we require. In order to accomplish this, we simply
determine the amount of solute needed, dissolve this in about 1/2 the amount of solvent and then
fill up to the proper amount of solution. It is best to use a volumetric flask for this purpose.
Example :
Describe how you would make 125 mL of a 0.250 M solution of sodium hydroxide.
Hebden p103 #95 - 98
Solutions wks #1
END OF LESSON 5
Ionic concentrations
When a solid dissolves in a solvent it can act one of three ways - it can either dissociate into
its ions (only if it is an ionic compound);it can remain as a molecule dissolved in water (if it is not an
ionic compound);or it can react with the water (we will not consider this at the present time). For the
sake of Chemistry 11, we will discuss only what happens when an ionic substance dissolves in a
solvent.
First of all, why do some ionic compounds dissolve readily in water while some others appear
to not be as readily soluble?
The reasons for this are due to the relative shapes and structures of both the solute and solvent
molecules. Sometimes the molecules have a slight charge on one side of the molecule compared to
the other. This type of molecule is called POLAR. Water is a good example of this. A NON-POLAR
molecule, such as carbon tetrachloride, can also be used as a solvent. The polarity of the molecules
is the basis of dissolution, but we will not get into that aspect of chemistry in this course.
When an ionic compound dissolves in water, it breaks into its ions.
Example :
NaCl (s) ----> Na+(aq) + Cl-(aq)
CaF2 (s) ----> Ca2+(aq) + 2 F-(aq)
We are going to assume that all of the compound breaks into its ion state. This is not always
true, but we will leave that for Chem 12.
If all of the compound ionizes, then we actually have none of the solute molecules left in the
solution. By looking at the ratios between the ions and compound (from the dissociation equation)
we can determine the concentrations of each individual ion in the solution.
(The abbreviation for concentration of an ion is to place the symbol for the ion inside square
brackets. For instance, shorthand for 'the concentration of the hydrogen ion' is [H+].)
Examples :
Find the concentration of each ion in the following solutions :
a) 12.0 M HCL
b) 18.0 M H2SO4
c) 1.15 M Cr2(SO4)3
Dilutions
When a solution of known concentration has its volume changed, it stands to reason that its
concentration changes also. If the liquid we are adding does not contain any of the original solute
OR if a reaction does NOT take place, the amount of dissolved solute does not change - only the
volume changes. In fact the concentration will change by a factor equal to the change in volume. In
short, the new concentration can be found by using the following formula :
C1V1=C2V2
new concentration = old concentration X
old volume
-------------------new volume
Examples :
1) 125 mL of 3.00 M HCl has 35.8 mL of water added to it. What is the new concentration ?
Solution :
2) 15.0 mL of a solution of 1.50 M HCl has some water added to it changing the concentration of
the acid to 1.15 M. What volume of water was added ?
Solution :
3) How much 6.00 M NaOH would we have to use to produce 275 mL of 0.500 M NaOH ?
Solution Worksheet #2
Mixing together solutions of different concentrations
As we said last class, if we mix together two solutions of different concentrations together,
the formula we used can NOT be used. In order to find the new concentration, we must find the
amount of solute in each solution.
Example : Determine the new concentration of HCl if 225 mL of 1.50 M HCl is mixed with 115 mL
of 2.50 M HCl.
Solutions Worksheet #3
END OF LESSON 6 – QUIZ #2 NEXT DAY
Reactions in Solutions
1. Precipitation Reaction:
Demo: Mix 0.2 M Pb(NO3)2 + 0.2 M KI. What ions are involved in this mixture?
The ions in solution are free to move around. If a reaction occurred, what ion combinations would be
possible?
The Chemical reaction:
The above reaction is a Double replacement reaction. The way the equation is written is called
CHEMICAL EQUATION.
On the product side of the equation we have a yellow solid, PbI2 (we'll figure out how to determine
which is the precipitate later) and a solution of KNO3. Rewriting the equation to show what we really
have in solution:
This is called an OVERALL IONIC EQUATION and is not used that often because of a number of
reasons - not the least of which is that it takes up too much space !!
If we look carefully at the overall ionic equation, we see that there are 2NO 3-(aq) ions and 2K+(aq) ions
on either side of the equation. If we cancel them out, we get
This is called a NET IONIC EQUATION for the reaction shown above. It shows exactly what is
happening in the solution. The NO3-(aq) ions and the K+(aq) ions are actually doing nothing in the
reaction. They can be thought of as just watching the lead ions and the iodide ions get together. For
this reason, they are referred to as SPECTATOR IONS. Spectator ions are never included in net
ionic equations.
Another example :
CuSO4 +
NaOH --->
- assume that the precipitate is Cu(OH)2
overall ionic equation:
net ionic equation:
Example : if we knew we got a precipitate of Ca3(PO4)2 in a reaction, the net ionic equation would be:
Examples: Write the a) chemical equation, b) overall ionic equation, and c) net ionic equation when
the following chemicals are mixed together:
a) NaOH + Fe(NO3)3 (solid formed is Fe(OH)3)
b) CaCl2 + Na2CO3 (solid formed is CaCO3)
c) AgNO3 + CaCl2 (solid formed is AgCl)
d) Na2SO4 + Ba(NO3)
(solid formed is BaSO4)
Solutions Worksheet #4
Reactions in Solutions
2. Acid Base Reaction:
This should be review for you!!! In short, the chemical properties of solutions of acids and
bases are as follows:
Acids :
Bases :
1) taste sour
2) conduct an electrical current
3) cause certain dyes to change colour
(e.g. litmus goes red)
4) liberates hydrogen when it reacts with
certain metals
5) loses the above properties when mixed
with a base though the resulting solution
conducts electricity.
1) taste bitter
2) conduct an electrical current
3) cause certain dyes to change colour
(e.g. litmus goes blue, phenolphthalein
goes pink)
4) feels slippery
5) loses the above properties when mixed
with a base though the resulting solution
conducts electricity.
6) INCREASES [H+]
6) INCREASES [OH-]
In their respective chemical formula, and more importantly when they dissolve in water, acids
have an H+ ion while bases have an OH- ion. (This is an incomplete definition of an acid and a base
but will suffice for chemistry 11.)
When an acid and a base react, the result is a salt and water. The type of reaction is called a
NEUTRALIZATION reaction because the properties of the acid and the base are neutralized by each
other. The net ionic reaction for this reaction is as follows :
H+(aq) + OH-(aq) ---> H2O (l)
We must remember that in any neutralization reaction, the moles of H+ MUST equal the moles
of OH-.
Example : What volume of 6.00 M HCl must be added to 125 mL of 1.59 M NaOH in order to
neutralize it ?
Example : What volume of 0.10 M Ca(OH)2 is needed to neutralize 115.2 mL of 0.55 M H3PO4?
Incomplete Neutralizations:


When acids and bases are mixed together and the moles of OH- and the moles of H+ are not
equal to one another.
Moles of OH- or moles of H+ are going to be left over in solution.
Example: Calculate the [OH-] and [H+] which results when 50.0 mL of 0.150 M NaOH is added to
50.0 mL of 0.200 M HCl.
Note that the concentration of OH- is not really equal to zero. But we will not deal with this concept
this year.
Example: Find the number of moles of HCl that must be added to 40.0 mL of 0.180M NaOH to
produce a solution having [OH-]=0.0316M.
Solutions Worksheet #5
END OF LESSON #7
Formation of precipitates

We will use the qualitative definition here that SOLUBLE means that we can make at least a
0.1 M solution of the compound at 25oC without producing as solid. How can we predict if a
precipitate will form ?
 The answer is not easy. In fact we don't have too many rules that always apply. Chemists have
done thousands of experiments and have collated the results into pages and pages of data tables.
 A table showing the solubility trends will be provided at first, but you will eventually need to know
the solubility trends. To use this table, we determine what the product will be in a reaction
(remember, the reactions here are double replacement reactions.) From here we must check out
both products to see if there is a precipitate formed.
Examples :
1) Identify the precipitate (if any) when the following solutions are mixed:
a) lead (II) nitrate and sodium chloride
b) ammonium hydroxide and copper (I) sulfate
c) potassium phosphate and cesium sulfide
2) Write the net ionic equation for the formation of any precipitate when solutions of the following are
mixed. If no precipitate is formed, write NO REACTION.
a) barium sulfide and ammonium carbonate
b) silver nitrate and sodium acetate
c) iron (III) chloride and copper (II) chloride
d) lead (II) acetate and lithium phosphate
e) Calcium nitrate and sodium hydroxide
f) lead (II) nitrate and potassium iodide
g) chromium (III) chloride + sodium hydroxide
3. When the following solutions are mixed together, what precipitate (if any) will form? If a reaction
occurs, write the net ionic equation.
a) barium chloride + sodium sulphate:
b) lead (II) nitrate + potassium chloride:
c) silver nitrate + sodium phosphate:
d) sodium hydroxide + ferric nitrate:
e) ferrous sulphate + potassium chloride:
f) aluminum nitrate + barium hydroxide:
g) calcium chloride + sodium sulphate:
h) potassium sulphide + nickel (II) nitrate:
4. Are the following compounds soluble or low in solubility?
a) Fe(OH)3
b) Hg2Cl2
c) PbSO4
d) BaCrO4
e) sodium sulphate
Bf) aluminum hydroxide
g) KNO3
h) NaNO3
i) FeSO4
k) CuCl2
Solutions Worksheet #6
END OF LESSON 8 – QUIZ #3 & Lab 16D Next Day
Learning Objectives for the Solutions Unit Test









Students should be able to define solute, solvent, solution, saturated, unsaturated and supersaturated and
solubility.
Students should be able to explain the super saturation demo and factors that affect solubility as well as be able
to read a solubility chart
Students should understand the connections between bond theories and the conductivity of ionic and covalent
aqueous solutions
Student should be able to identify polar and non-polar solvents and use this to predict solubility of give solutes
Students should be able to use lewis structures in conjunction with VSEPR to predict the shape and bond angles
for molecules with up to 4 central bonds
Students should understand what resonance structures are
Students should recongize and be able to explain the differences between dissociation and ionization
Students must be able to calculate molarity, dilutions, concentration of ions in solutions, and mixing of solutions
as well as stoichiometry involving titrations.
Students should be able to write overall and net ionic equations, identify spectator ions and predict the
formation of precipitates using a solubility chart.
Review Worksheet #1
Review Worksheet #2