Chemistry 30
Unit 5:
Solutions and
Solubility
Components of a Solution

A solution is composed of two or more _________________________ substances mixed together, one
of which is a solvent and the other a solute.

A cup of instant coffee is the _________________________, the hot water is the _________________________,
and the instant coffee is the _________________________.

A _________________________ is a solution that contains the maximum amount of solute dissolved in
a solvent. A __________________________________________________ is a solution where conditions have
been changed to allow more solute to dissolve than would at room temperature (this is done
by heating the solvent or solution).

When the solvent and solutes are liquids we can use the terms miscible and immiscible.
_________________________ means that the two liquids will mix together, and _________________________
means that the two liquids will not dissolve in one another.

Remember, the solvent and solute can be solid, liquid OR gas. There are
______________________________for solutions where the solute and solvent are both solids and when
the solute is a liquid, and the solvent is a solid.
o Solid-in-solid – _________________________ (e.g., brass = zinc + copper, sterling silver = silver
+ copper).
o Liquid-in-solid – _________________________ (e.g., older dental fillings = mercury + silver).

Factors that may affect solubility are temperature and pressure:
o Temperature: Gases: Temp ⬆, solubility ⬇
Solids and liquids: Temp ⬆, solubility ⬆
o Pressure: Gases: Pressure ⬆, solubility ⬆
Solids and liquids: unaffected

Like dissolves like. That is polar covalent molecules usually dissolve in other polar covalent
molecules and non-polar covalent molecules dissolve in non-polar covalent molecules.
RECALL: Polarity can be determined by subtracting electronegativities. A polar molecule is one where one (or more)
atom in the molecule has a stronger pull on the electrons than another atom, so the electrons move closer to the
atom with the stronger pull (higher electronegativity). This creates a dipole, where the atom with the stronger pull
has a slightly negative charge and the atom with the lower electronegativity (less pull) has a slightly positive charge.
Water as a Solvent:

Solutions in which water is the
solvent are called
_________________________ solutions.

Most covalent liquids are only
able to dissolve other covalent
compounds; water can dissolve
both _________________________and
covalent compounds.
The attraction of water
_________________________ for ions pulls ions out of a crystalline lattice and into aqueous solution.
Aqueous Solutions:

Sometimes when we mix two dissolved salts together, we end up with a _________________________.
This is because a new ionic compound is formed that is insoluble in water.
We can use solubility charts or solubility tables to determine if a solid will be formed.
Ion
NO3–
ClO4–
Cl–
I–
BrSO42CO32PO43OH–
S2Na+
K+
NH4+
Solubility
soluble
soluble
soluble
soluble
soluble
soluble
insoluble
insoluble
insoluble
insoluble
soluble
soluble
soluble
Exceptions
none
none
except Ag+, Hg22+, *Pb2+
except Ag+, Hg22+, *Pb2+
except Ag+, Hg22+, *Pb2+
except Ca2+, Ba2+, Sr2+, Hg2+, Pb2+, Ag+
except Group IA and NH4+
except Group IA and NH4+
except Group IA, *Ca2+, Ba2+, Sr2+
except Group IA, IIA and NH4+
None
None
None
* = slightly soluble
Ex. If we mix sodium ions with hydroxide ions will a precipitate form?
Ex. If we dissolved sodium hydroxide in calcium chloride, will a precipitate form?
Ex. Calcium nitrate with sodium carbonate
See Components of Solutions Assign
Concentration

Concentration refers to the amount of solute dissolved in a specific _________________________of
solvent.

Concentrated and dilute are _________________________ terms we use to describe concentration.
Dilute means there is not a lot of solute in the solution and concentrated means there is a lot of
solute in the solution.

Common concentrations are molarity, molality, ppm, and ppb.
Molarity
This is the number of moles of solute dissolved in one litre of solution.
The formula for molarity is as follows:
n = # moles of solute
V = volume of solution in litres
c = concentration in moles per litre (M)
Ex. If a teaspoon (5.0mL) of a 0.50 M solution of NaCl was evaporated, how many moles of sodium
chloride would be left? What mass of NaCl would be left?
Ex. Antifreeze is a solution of ethylene glycol, C2H6O2, in water. If 4.50 L of antifreeze contains 2.00 kg of
ethylene glycol, what is the concentration of the solution?
Standard solution:

A standard solution is a solution of a known concentration. This means it has a precise
_________________________ of solute in a specific _________________________ of solution.

_________________________ solutions are used in experiments where the concentration of a solution
must be known.

We use volumetric flasks to prepare standard solutions as they have a small margin of
_________________________ when compared to other pieces of lab equipment. Volumetric flasks come
in a variety of sizes (volumes).
To prepare a standard solution we follow the following steps:
1. Calculate the required _________________________ of solute needed using the volume and
concentration you want to end up with.
2. Weigh out the mass of the solute needed and add it to a volumetric flask of the appropriate size.
3. Dissolve the solid in _________________________ water using less than _________________________
of the final solution volume.
4. Once the solute is dissolved, add the rest of the water. Be sure to use a medicine
dropper for the final few milliliters of water. Use the calibration line to set the
_________________________ in the appropriate spot.
Ex. Describe the preparation of 2.000 L of a standard aqueous solution containing 0.1000 mol/L
potassium nitrate.
To prepare the solution we place_______________ g of KNO3 in a 2.000 L volumetric flask._________________________________
the required amount of water is added. When all the potassium nitrate is dissolved, the solution is diluted with the
remaining water.
Dilution Calculations:


When making a solution in chemistry laboratories you usually only have access to substances
with high concentration solutions (called stock solutions) and are then required to dilute the
stock solutions. A calculation needs to be completed in order to determine the amount of
distilled water that needs to be added to a certain volume of stock solution in order to create the
desired concentration.
Since the number of _________________________ of solute in a solution does not change when you
dilute it, the equation for dilution is as follows:
Ex. Water is added to 200.mL of 2.40M ammonia cleaning solution (NH3), until the final volume is 1.00L.
Find the molar concentration of the final diluted solution.
Ex . What volume of concentrated sulphuric acid (containing 18.0 M H2SO4) is required to prepare 5.00
L of 0.150 M aqueous sulphuric acid solution by dilution with water?
Ion Concentration:
Consider a 0.20M aqueous solution of sodium carbonate. The sodium carbonate will be completely
dissociated into ions:
The concentration of the sodium ions can be calculated using the conversion factors from the balanced
equation (similar to mol- mol stoichiometry). Remember: concentration is represented by using square
brackets.
Ex. What are the concentrations of the ions in an aqueous solution containing 0.15 mol/L iron(III)
nitrate?
Ex. 250mL of 0.30M K2SO4 and 250mL of 0.80M MgCl2 are mixed and no reaction results. What is the
concentration of each substance in the final solution, and the concentration of each individual ion?
Ppm & ppb:

Usually in chemistry we work with _________________________ when discussing concentration. When
working with very dilute solutions (solutions with low concentrations) it is more convenient to
work with ________________ or ________________.

A “__________________________” refers to a gram.

A concentration of 1ppb is equivalent to 1g of solute dissolved in ________________________ of solvent.

The allowable concentration of ________________ chemicals is usually measured in ppm or ppb.

For example the allowable concentration of mercury in air is _____________________, this means there
are ___________________________________. The allowable concentration of copper in drinking water is
1ppm, this means ________________________________ .

To do ppm or ppb calculations we assume that 1g of water has a volume of _______________.
Ex. Change 2.00M KNO3 to ppm.
Ex. The concentration of gold in the ocean is 0.005ppb. What would be the molarity of gold in a sample
of ocean water.
See Concentration and Dilution Assignment
The Solubility Product Constant:

All salts have different solubilities. Even ‘______________________________’ salts will dissolve in water
to some extent.

When a salt is dissolved in water, some of the salt dissolves and some may not. The amount of
salt that dissociates into ions is _____________________________ for every salt.

Ksp is a number that represents how much of the salt will _____________________________ into ions.

Ksp is equal to the product of the concentration terms each raised to the power of the coefficients
of the substance in the dissociation equation.

Ksp values @ 25oC can be found on the table on the next page.

The equilibrium constant Ksp called the _____________________________ indicates how soluble a
substance is.
o If the Ksp is low, the solid is ____________________________________________.
o If the Ksp is high, the solid is _____________________________.
An expression for the solubility product constant, Ksp, of calcium phosphate would be as follows:
Balanced Equation:
Ksp is equal to the product of the concentrations of the products, each raised to the power of the
coefficient in the balanced equation, over the concentration of the reactant to the power of the
coefficient.
When using Ksp we do not include the concentrations of solids, liquids or gases, as they are not
dissolved. Therefore, the concentration of the reactant should not be included, as it is a solid.
Example: Calculate the solubility product constant for calcium fluoride at 20oC given that it`s solubility
is 1.6 × 10−2 g/L.
First write the chemical equation:
Then convert to molarity (the number of moles of CaF2 that dissolved in 1L):
Find Ksp using solubility constant expression:
Solubility Product Constants near 25 oC
Formula
Al(OH)3
AlPO4
BaCO3
BaCrO4
BaF2
Ba(OH)2
BaSO4
BaSO3
BaS2O3
BiOCl
BiOOH
CdCO3
Cd(OH)2
CdC2O4
CdS
CaCO3
CaCrO4
CaF2
CaHPO4
Ca(OH)2
CaC2O4
Ca3(PO4)2
CaSO4
CaSO3
Cr(OH)2
Cr(OH)3
CoCO3
Co(OH)2
Co(OH)3
CoS
CuCl
CuCN
CuI
Cu3(AsO4)2
CuCO3
CuCrO4
Cu[Fe(CN)6]
Ksp
1.3×10–33
6.3×10–19
5.1×10–9
1.2×10–10
1.0×10–6
5×10–3
1.1×10–10
8×10–7
1.6×10–6
1.8×10–31
4×10–10
5.2×10–12
2.5×10–14
1.5×10–8
8×10–28
2.8×10–9
7.1×10–4
5.3×10–9
1×10–7
5.5×10–6
2.7×10–9
2.0×10–29
9.1×10–6
6.8×10–8
2×10–16
6.3×10–31
1.4×10–13
1.6×10–15
1.6×10–44
4×10–21
1.2×10–6
3.2×10–20
1.1×10–12
7.6×10–36
1.4×10–10
3.6×10–6
1.3×10–16
Formula
Cu(OH)2
CuS
FeCO3
Fe(OH)2
FeS
FeAsO4
Fe4[Fe(CN)6]3
Fe(OH)3
FePO4
Pb3(AsO4)2
Pb(N3)2
PbBr2
PbCO3
PbCl2
PbCrO4
PbF2
Pb(OH)2
PbI2
PbSO4
PbS
Li2CO3
LiF
Li3PO4
MgNH4PO4
Mg3(AsO4)2
MgCO3
MgF2
Mg(OH)2
MgC2O4
Mg3(PO4)2
MnCO3
Mn(OH)2
MnS
Hg2Br2
Hg2Cl2
Hg2I2
HgS
Ksp
2.2×10–20
6×10–37
3.2×10–11
8.0×10–16
6×10–19
5.7×10–21
3.3×10–41
4×10–38
1.3×10–22
4×10–36
2.5×10–9
4.0×10–5
7.4×10–14
1.6×10–5
2.8×10–13
2.7×10–8
1.2×10–15
7.1×10–9
1.6×10–8
3×10–28
2.5×10–2
3.8×10–3
3.2×10–9
2.5×10–13
2×10–20
3.5×10–8
3.7×10–8
1.8×10–11
8.5×10–5
1×10–25
1.8×10–11
1.9×10–13
3×10–14
5.6×10–23
1.3×10–18
4.5×10–29
2×10–53
Formula
NiCO3
Ni(OH)2
NiS
ScF3
Sc(OH)3
AgC2H3O2
Ag3AsO4
AgN3
AgBr
AgCl
Ag2CrO4
AgCN
AgIO3
AgI
AgNO2
Ag2SO4
Ag2S
Ag2SO3
AgSCN
SrCO3
SrCrO4
SrF2
SrSO4
TlBr
TlCl
TlI
Tl(OH)3
Sn(OH)2
SnS
ZnCO3
Zn(OH)2
ZnC2O4
Zn3(PO4)2
ZnS
* Sulfide equilibrium of the type:
MS(s) + H2O(l) ⇌ M2+(aq) + HS–(aq) + OH–(aq)
Ksp
6.6×10–9
2.0×10–15
3×10–19
4.2×10–18
8.0×10–31
2.0×10–3
1.0×10–22
2.8×10–9
5.0×10–13
1.8×10–10
1.1×10–12
1.2×10–16
3.0×10–8
8.5×10–17
6.0×10–4
1.4×10–5
6×10–51
1.5×10–14
1.0×10–12
1.1×10–10
2.2×10–5
2.5×10–9
3.2×10–7
3.4×10–6
1.7×10–4
6.5×10–8
6.3×10–46
1.4×10–28
1×10–26
1.4×10–11
1.2×10–17
2.7×10–8
9.0×10–33
2×10–25
Determining Individual Ion Concentration from Ksp:
Ex. Calculate the concentrations of barium ions and sulphate ions in a saturated aqueous solution of
barium sulphate at 25oC.
Ex. When 10.0 L of a saturated solution of magnesium carbonate, at 22°C is evaporated to dryness,
0.12g of solid magnesium carbonate is obtained. Calculate the Ksp of magnesium carbonate.
Ex. A saturated solution of ionic calcium hydroxide has a hydroxide ion concentration of 3.0×10-3mol/L.
Calculate the Ksp of calcium hydroxide at the same temperature.
See Solubility Product Constant Assignment
The Common Ion Effect:

Remember that when a salt dissolves in water, some of it dissolves and some may stay as a solid.
There is a point at which the salt dissolving reaches _____________________________. This means that
for each formula unit of salt that dissolves, a pair of ions come together to create a formula unit.

When we dissolve ions in water we know that the product of the concentrations of the ions, to
the power of their coefficients in a balanced equation, is equal to the ____________.

As long as the product of the concentration of the ions does not exceed the Ksp value, no
precipitate forms.
Let’s look at the following example:
Copper (I) iodide is not very soluble in water (Ksp = 4 × 10−19): CuI(s) ⇌Cu+(aq) + I−(aq)
Since Ksp = [Cu+][I-], if the product of the concentrations of Cu+ and I- stays below 4 × 10−19, no
precipitate will form, however, if we add extra iodine or copper ions (the common ion) things
will change. Let`s say we added some iodide ions; this will increase the concentration of
_____________________________ and the product of the copper ions and iodide ions would be
_____________________________ than the Ksp value. This is not possible so the concentrations would
need to be lowered. In order to lower the concentration of iodide ions, some copper (I) iodide
would need to _____________________________ out. This follows Le Chatelier`s Principle because we
increased the concentration of the products (iodide ions) so the reaction shifts to the
_____________________________.
After adding iodide ions to the solution, copper (I) iodide will precipitate out, but once the
reaction reaches equilibrium again, the concentration of copper ions and iodide will no longer
be _____________________________ (there will be more iodide ions in solution than copper ions).
This also means the concentration of Cu+ will be lower than at the original equilibrium and the
concentration of I- will be higher than at the original equilibrium.
Overview: If a slightly soluble ionic compound is dissolved in water, you can force
precipitation of that salt by adding a readily soluble ionic compound that has an ion in
common with the slightly soluble salt. This shift is known as the common ion effect.
Ex. The Ksp of silver iodide is 8.3 1017 . What is the iodide ion concentration of a 1.00 L saturated
solution of AgI to which 0.020 mol of AgNO3 is added?
AgI(aq) ⇌ Ag+(aq) + I-(aq)
Ksp = [ Ag  ][ I  ]  8.3 1017
Selective Precipitation

Selective precipitation is when you separate ions from two _____________________________ to create a
precipitate.

To calculate selective precipitation we use _____________________________ concentrations (Q), instead of
the equilibrium concentrations.

In selective precipitation there are restraints in forming a precipitate; a precipitate doesn't always
form in a solution. It depends on the relationship between K and Q:
o If Q>K _________precipitate forms and the reaction proceeds to the left
o If Q<K _________precipitate forms and the reaction proceeds to the right
o If Q=K the reaction is at _____________________________
Example: If 40L of 0.0050M KCl is mixed with 60L of 0.0030M Pb(NO3)2. Will the precipitate PbCl2
(Ksp= 1.6 X 10-5) form?
First we calculate the initial concentrations of Pb2+ and Cl-:
Calculate value of Q:
Example: If 100L of 0.03M Pb(NO3)2 is mixed with 200L of 0.09M KCl. Does a precipitate (PbCl2) form?
(Ksp = 1.6 X 10-5). What are the equilibrium concentrations of Pb2+ and Cl-?
By using selective precipitation, we can calculate the amount of a reagent necessary to react with the
salt.
Example: If a solution contains 1.0 x 10-5 M of Pb2+ and 2.0 x 10-5 M Ag+ and Cl- is slowly added to
the solution, will AgCl (Ksp=1.6 x 10-10) or PbCl2 (Ksp=1.6 x 10-5) precipitate first? Which will be in
excess?
See Common Ion and Selective Precipitation Assignment
Lab: Analyzing for Anions
PURPOSE
In this experiment you will observe and record some characteristic chemical reactions of several
negatively charged ions: SO42-, CO32-, and I- using selected reagents. From a study of these observations
you will be able to develop a method for identifying each of these anions in the presence of the other
ions.
Your work will be typical of the work done by chemists who have developed elaborate schemes which
can be used to identify dozens of cations and anions in a given sample. Organize a data table in which
you can record your observations for the results obtained by adding each of the reagents to each of the
solutions of anions to be studied.
MATERIALS
Anion obtained
Cation obtained
Na2SO4, sodium sulfate, 0.1 M SO42-
Ba(NO3)2 , barium nitrate, 0.1 M Ba2+
Na2CO3, sodium carbonate, 0.1 M CO32-
AgNO3, silver nitrate, 0.1 M Ag+
NaCI, sodium chloride, 0.1 M CI-
HNO3, nitric acid, 1.0 M H+
Nal, sodium iodide, 0.1 M IPROCEDURE
a. The tests can easily be made by putting a few drops of a solution on a plastic film and adding to this a
few drops of one of the reagents. DO NOT use the tip of a dropper from a stock bottle for stirring. Doing
so could contaminate the solution contained in the bottle. Use a thin glass rod.
b. Test 2-3 drops of each anion solution separately with 1-2 drops of the 0.1 M Ba(NO3)2 reagent.
Record the results observed. To each mixture that contains a precipitate add 2-3 drops of 1.0 M HNO3
and record any changes you observe.
c. Test 2-3 drops of each anion solution separately with 1-2 drops of the 0.1 M AgNO3 reagent. Record
the results observed. To each mixture that contains a precipitate add 2—3 drops of 1.0 M HNO3 and
record any changes you observe.
d. Prepare new samples of the silver precipitates which formed in step c. Add 2—3 drops of 6 M
aqueous ammonia, NH3(aq), to the precipitates and record any changes you observe.
e. Test 2—3 drops of each anion solution separately with 1—2 drops of 1.0 M HNO3 and record the
results.
f. Examine the pattern of reactions shown by your table of observations. Develop a scheme by which
you could identify either one or a combination of two of the anions in a single solution. Organize your
scheme in such a way that another student could follow it. You may wish to try it out by making up your
own "trial unknowns" containing various combinations of the anions.
g. Obtain an unknown from your teacher and analyze it for the four anions. Report your results. Include
supporting evidence for your conclusions.
ANALYSIS
1) Write net ionic equations for the precipitation reactions which occurred when solutions of the anions
were mixed with the reagent containing Ba2+ ions.
2) Do the same for the precipitation reactions which occurred in step c with the Ag+ ions.
3) Write equations for the reactions which occurred when the precipitates were acidified with 1.0 M
HNO3 .
4) Write equations for the reactions which occurred when the silver precipitate reacted with aqueous
ammonia to form the complex ion, Ag(NH3)2+.
Name:_______________________________________
Components of Solutions and Precipitate Formation
1. Cola drinks contain water, sugar, flavouring, and colouring agents, and carbon dioxide gas under
pressure. Identify the solutes and solvent in this solution.
2. Write chemical formulas for the following compounds, and using your solubility rules, predict
their solubilities in water.
a)
potassium phosphate
b)
calcium carbonate
c)
copper(II) bromide
d)
aluminum sulfide
e)
nickel(II) sulfate
3. A solution contains both Cu2+(aq) and Ag+(aq). To precipitate one of these ions and leave the other
in solution, which of the following should be added?
a)
H2S(aq)
c)
HNO3(aq)
b)
HCl(aq)
d)
NH4NO3(aq)
4. Use your Solubility Rules to predict whether a precipitate will form when aqueous solutions of
the following substances are mixed. If a precipitate is expected, give the balanced net ionic
equation.
a) sodium sulfide and chromium(III) sulfate
b) sodium sulfate and iron(III) nitrate
c) potassium hydroxide and cobalt(II) chloride
d) aluminum sulfate and sodium carbonate
5. Write balanced equations that show which ions are produced when the following ionic
substances are dissolved in water:
a)
sodium sulfide
b)
calcium iodide
c)
lithium carbonate
d)
iron(III) sulfate
e)
cobalt(II) nitrate
6. To an aqueous solution containing the cations Ca2+, Ag+, Cu2+, and K+ is added a sodium bromide
solution. The precipitate A is filtered off and a sodium sulfide solution is added to the filtrate.
The black precipitate B that forms is removed by filtration. A solution of sodium carbonate is
added next, again resulting in a precipitate C. Identify A, B, and C, and give the balanced net ionic
reactions for their formation.
7. An aqueous solution may contain any or all of the cations Sr2+, Zn2+, and Hg22+. Aqueous solutions
of sodium sulfide, sodium chloride and sodium carbonate, may be added to determine which
cations are present. Write the balanced net ionic equations, and provide a flowchart to illustrate
the order in which the test solutions should be used.
Name:____________________________
Concentration and Dilution
1. What happens to the concentration of a solution when
a)
more solute is added?
b)
the solution is heated and the solvent evaporates?
c)
part of the solution is poured off?
2. Explain how you would prepare 500.0mL of a 0.100M solution of magnesium sulfate from solid
compound?
3. Calculate the concentrations of the four most common ions in sea water.
Major Ions Present in Sea Water
Ion
Concentration (g/L) % of sea salt
−
Cl
18.8
53.7
Na+
10.8
30.8
2−
SO4
2.65
7.6
2+
Mg
1.29
3.7
Ca2+
0.412
1.2
+
K
0.38
1.1
−
HCO3
0.14
0.4
Br−
0.067
0.2
4. Determine the mass of solute present in the following solutions:
a)
a 2.00 L solution of 2.00 mol/L sodium hydroxide
b)
a 50.00 mL aqueous solution containing 4.550 × 10-3 M copper(II) nitrate
5. Calculate the concentrations of the following solutions:
a)
a 250.0 mL aqueous solution containing 25.50 g of sodium phosphate
b)
a 75.00 mL solution of concentrated aqueous ammonia (14.8 M), diluted with
water to a volume of 2.000 L.
6. What mass of lead(II) nitrate is present in 25.00 mL of a 0.1000 M aqueous solution?
7. Calculate the concentrations of the following solutions:
a)
a 0.750 L aqueous solution containing 90.0 g ethanol, C2H5OH
b)
a 25.00 mL solution of 0.1250 M iron(II) sulfate, diluted with water to a volume of 500.0
mL
8. What volume of concentrated hydrochloric acid, containing 11.9 mol/L hydrogen chloride, is
needed to make 2.00 L of 0.200 M solution by dilution with water?
9. Find [HCl] if 25.0mL concentrated HCl (12.0M) is diluted to a final volume of 2.00L.
10. How would you make 1.00L of a 0.50M NH3 solution given 15M NH3 solution?
11. If 25mL of a 0.2M Li3PO4 solution was poured into 130mL of pure water, calculate [Li+].
12. Calculate the concentrations of ions in the following solutions:
a)0.015 M solution of sodium sulfate
b)2.00 L aqueous solution containing 17.1 g aluminum sulfate
13. What are the concentrations of the ions in the following solutions:
a) an aqueous solution containing 0.075 M strontium nitrate?
b)
a 200 mL aqueous solution containing 4.26 g sodium sulfate?
Name:_____________________________
Solubility Product Constant Assignment
1. Write the solubility product constant expression for equilibrium in a saturated solution of
a)
iron(III) hydroxide
b)
gold(III) oxalate
c)
magnesium fluoride
2. Which compound in each set is more soluble? How do you know?
a)
BaSO4 (Ksp = 1.1 × 10−10) or PbSO4 (Ksp = 1.6 × 10−8)
b)
PbI2 (Ksp = 7.1 × 10−9) or CaF2 (Ksp = 5.3 × 10−9)
3. What are the concentrations of the resulting ions in saturated aqueous solutions of the following
compounds?
a) lead(II) sulfate
Ksp = 1.30 × 10-8
b) silver carbonate
Ksp = 6.2 × 10-12
c) magnesium hydroxide
Ksp = 1.2 × 10-11
4. Which of the following is most likely to occur when 1.058 g of Na2SO4 is added to 375.0 mL of
saturated BaSO4(aq)?
a)
[Ba2+] is reduced.
c)
The solubility of BaSO4 is increased.
2−
b)
[SO4 ] is reduced.
d)
[Ba2+] does not change.
5. At 20°C, a saturated aqueous solution of silver carbonate contains 32 mg per liter of solution.
Calculate Ksp at 20°C.
6. Calculate the concentrations of ions in a saturated aqueous solution of the following:
a) silver iodide; Ksp = 1.5 × 10−16
b) strontium carbonate; Ksp = 1.6 × 10−9
7. What are the concentrations or the ions in saturated aqueous solutions of the following:
a) lead(II) iodide; Ksp = 1.4 × 10−8
b) silver sulfate; Ksp = 1.5 × 10−4
c) iron(III) hydroxide; Ksp = 6.3 × 10−38
8. Calculate the Ksp values of the substances from their solubility’s in water:
a) thallium(I) chloride; 3.4 g/L at 25°C
b) silver bromide; 1.3 × 10−4 g/L at 20°C
c) CaF2, solubility 3.32  10−4 mol/L
d) Aluminum hydroxide, solubility 2.6  10−9 mol/L
e) Cadmium iodate, solubility 0.097 g Cd(IO3)2/100 mL H2O
Name:______________________________
Common Ion Effect and Selective Precipitation
1. Will a precipitate of calcium fluoride form when 0.084 g of sodium fluoride is dissolved in 1.00 L
of a 0.010 M aqueous solution of calcium chloride? The Ksp for calcium fluoride is 3.9 × 10−11.
2. Would we expect a precipitate of silver bromate (Ksp = 5.9 × 10-5), to form when 50.00 mL of
0.0020 M aqueous solution of silver nitrate, is added to 250.0 mL of 0.020 M aqueous solution of
potassium bromate?
3. Will a precipitate of magnesium hydroxide (Ksp = 1.2 × 10-11) form when 1.00 mL of 0.010 M
aqueous solution of calcium hydroxide is added to 1.0 L of 0.20 M aqueous solution of
magnesium nitrate? The change in volume caused by the addition of 1.00 mL can be ignored.
4.
A 0.10 M aqueous solution of sodium sulfate is added one drop at a time to 1.00 L of a 0.0010 M
aqueous solution of lead(II) nitrate. What is the minimum volume that has to be added before a
precipitate of lead(II) sulfate starts to form? The Ksp for lead(II) sulfate is 1.3 × 10−8, and the
change in volume can be ignored.
5. Upon addition of hydroxide ions to sea water, magnesium hydroxide will precipitate. If the
magnesium ion concentration in sea water is 5.3 × 10−2 M, calculate the maximum hydroxide ion
concentration in sea water. The Ksp for magnesium hydroxide is 1.2 × 10−11.
6. A sample of sea water contains 5.3 × 10−1 M chloride ion and 8.4 × 10−4 M bromide ion. What
concentration of added silver ion would cause precipitation of silver chloride and silver
bromide? Which of these two silver halides would precipitate first? The Ksp for silver chloride is
1.6 × 10−10; the Ksp for silver bromide is 6.5 × 10−13.
7. What volume of 0.10 M aqueous calcium bromide must be added to 100.0 mL of 0.100 M
aqueous lead(II) nitrate before a precipitate of lead(II) bromide starts to form? Assume in your
calculation that the total volume remains 100.0 mL. The Ksp for lead(II) bromide is 1.4  10−8.
8. @25oC will a precipitate form, and if so, what is the precipitate when 235 mL of 0.0022 M MgCl2
and 485 mL of 0.0055 M NaF are mixed?
9.
If 1.00 mg of Na2CrO4 is added to 225 mL of 0.00015 M AgNO3(aq), will a precipitate form?
1.1 × 10−12
Ksp =
10. A 100.0 mL solution of 1.00 M aqueous sodium sulfate is added to a 200.0 mL solution of 1.00 M
aqueous barium chloride. Barium sulfate precipitates from the mixture. (The Ksp for barium
sulfate is 1.1  10−10.) Calculate the following:
a) mass of barium sulfate that precipitates. 1st Determine if a precipitate will form, if so
assume all of Ba2+ or all SO42- reacts (whichever is limiting).
b)
concentrations of all ions that remain in solution