Chemistry 30

Aqueous

Equilibria

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).

· Factors that may affect solubility are temperature and pressure:

· Temperature: Gases: Temp ⬆, solubility ⬇

Solids and liquids: Temp ⬆, solubility ⬆

· 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

Solubility

Exceptions

NO3–

soluble

none

ClO4–

soluble

none

Cl–

soluble

 except Ag+, Hg22+, *Pb2+

I–

soluble

except Ag+, Hg22+, *Pb2+

Br-

soluble

except Ag+, Hg22+, *Pb2+

SO42-

soluble

except Ca2+, Ba2+, Sr2+, Hg2+, Pb2+, Ag+

CO32-

insoluble

except Group IA and NH4+

PO43-

insoluble

except Group IA and NH4+

OH–

insoluble

except Group IA, *Ca2+, Ba2+, Sr2+

S2-

insoluble

except Group IA, IIA and NH4+

Na+

soluble

None

K+

soluble

None

NH4+

soluble

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. If we mix aqueous solutions of calcium nitrate with an aqueous solution of sodium carbonate will a precipitate form?

NOTE: Compounds that are insoluble in water form solids (represented by subscripted (s) in balanced equations) and compounds that are soluble in water are aqueous (represented by subscripted (aq)).

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?

Solubility Curves:

There are charts and tables available that we can use to get an idea of how soluble a certain solute is in a certain solvent. Solubility ___________________ tell us what mass of solute will dissolve in 100g of water over a range of temperatures.

You'll notice that for most substances, solubility _____________________ as temperature increases. In solutions involving liquids and solids typically more solute can be dissolved at higher temperatures.

Ex. At 30°C approximately __________ of KClO3 will dissolve in 100g of water. If the temperature is increased to 80°C, approximately ____________ of the substance will dissolve in 100g (or 100mL) of water.

Ex. 2 What mass of solute will dissolve in 100mL of water at the following temperatures. Also determine which of the three substances is most soluble in water at 15°C.

1.

KNO3at 70°C

 

2.

NaCl at 100°C

 

3.

NH4Cl at 90°C

 

· On a solubility curve, the lines indicate the concentration of a _____________________ solution - the ________________________ amount of solute that will dissolve at that specific temperature. The molar concentration of the substance can be calculated.

· The values on the graph ________________________ a curve represent unsaturated solutions - more solute could be dissolved at that temperature. Values ________________________a curve represent supersaturated solutions, a solution which holds more solute that can normally dissolve in that volume of solvent.

Ex 3. Determine the molarity of a saturated NaCl solution at 25°C.

Ex 4. What term - saturated, unsaturated, or supersaturated - best describes:

· a solution that contains 70g of NaNO3 per 100 mL H2O at 30°C

· a solution that contains 60g of dissolved KCl per 100 mL H2O at 80°C

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?

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.

· If the Ksp is low, the solid is ____________________________________________.

· 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

Ksp

Al(OH)3

1.3×10–33

AlPO4

6.3×10–19

BaCO3

5.1×10–9

BaCrO4

1.2×10–10

BaF2

1.0×10–6

Ba(OH)2

5×10–3

BaSO4

1.1×10–10

BaSO3

8×10–7

BaS2O3

1.6×10–6

BiOCl

1.8×10–31

BiOOH

4×10–10

CdCO3

5.2×10–12

Cd(OH)2

2.5×10–14

CdC2O4

1.5×10–8

CdS

8×10–28

CaCO3

2.8×10–9

CaCrO4

7.1×10–4

CaF2

5.3×10–9

CaHPO4

1×10–7

Ca(OH)2

5.5×10–6

CaC2O4

2.7×10–9

Ca3(PO4)2

2.0×10–29

CaSO4

9.1×10–6

CaSO3

6.8×10–8

Cr(OH)2

2×10–16

Cr(OH)3

6.3×10–31

CoCO3

1.4×10–13

Co(OH)2

1.6×10–15

Co(OH)3

1.6×10–44

CoS

4×10–21

CuCl

1.2×10–6

CuCN

3.2×10–20

CuI

1.1×10–12

Cu3(AsO4)2

7.6×10–36

CuCO3

1.4×10–10

CuCrO4

3.6×10–6

Cu[Fe(CN)6]

1.3×10–16

Formula

Ksp

Cu(OH)2

2.2×10–20

CuS

6×10–37

FeCO3

3.2×10–11

Fe(OH)2

8.0×10–16

FeS

6×10–19

FeAsO4

5.7×10–21

Fe4[Fe(CN)6]3

3.3×10–41

Fe(OH)3

4×10–38

FePO4

1.3×10–22

Pb3(AsO4)2

4×10–36

Pb(N3)2

2.5×10–9

PbBr2

4.0×10–5

PbCO3

7.4×10–14

PbCl2

1.6×10–5

PbCrO4

2.8×10–13

PbF2

2.7×10–8

Pb(OH)2

1.2×10–15

PbI2

7.1×10–9

PbSO4

1.6×10–8

PbS

3×10–28

Li2CO3

2.5×10–2

LiF

3.8×10–3

Li3PO4

3.2×10–9

MgNH4PO4

2.5×10–13

Mg3(AsO4)2

2×10–20

MgCO3

3.5×10–8

MgF2

3.7×10–8

Mg(OH)2

1.8×10–11

MgC2O4

8.5×10–5

Mg3(PO4)2

1×10–25

MnCO3

1.8×10–11

Mn(OH)2

1.9×10–13

MnS

3×10–14

Hg2Br2

5.6×10–23

Hg2Cl2

1.3×10–18

Hg2I2

4.5×10–29

HgS

2×10–53

Formula

Ksp

NiCO3

6.6×10–9

Ni(OH)2

2.0×10–15

NiS

3×10–19

ScF3

4.2×10–18

Sc(OH)3

8.0×10–31

AgC2H3O2

2.0×10–3

Ag3AsO4

1.0×10–22

AgN3

2.8×10–9

AgBr

5.0×10–13

AgCl

1.8×10–10

Ag2CrO4

1.1×10–12

AgCN

1.2×10–16

AgIO3

3.0×10–8

AgI

8.5×10–17

AgNO2

6.0×10–4

Ag2SO4

1.4×10–5

Ag2S

6×10–51

Ag2SO3

1.5×10–14

AgSCN

1.0×10–12

SrCO3

1.1×10–10

SrCrO4

2.2×10–5

SrF2

2.5×10–9

SrSO4

3.2×10–7

TlBr

3.4×10–6

TlCl

1.7×10–4

TlI

6.5×10–8

Tl(OH)3

6.3×10–46

Sn(OH)2

1.4×10–28

SnS

1×10–26

ZnCO3

1.4×10–11

Zn(OH)2

1.2×10–17

ZnC2O4

2.7×10–8

Zn3(PO4)2

9.0×10–33

ZnS

2×10–25

* Sulfide equilibrium of the type:MS(s)  +  H2O(l)  ⇌  M2+(aq)  +  HS–(aq)  +  OH–(aq)

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 . 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 =

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:

· If Q>K  _________precipitate forms and the reaction proceeds to the left

· If Q

· 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 concentrations of Pb2+ and Cl- after the solutions have been mixed:

Calculate value of Q:

Example: If 1.00L of 0.03M Pb(NO3)2 is mixed with 2.00L of 0.09M KCl. Does a precipitate (PbCl2) form? (Ksp = 1.6 X 10-5).

By using selective precipitation, we can calculate the amount of a reagent necessary to react with the salt.

Example: If a solution contains  2.0 x 10-5 M  of  Pb2+  and 1.0 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?

See Common Ion and Selective Precipitation Assignment

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.

Solubility Curves Assignment

Use the Solubility Curve Provided for the following questions:

1. Describe how the solubility of solids (solid line) and gases (dashed line) is different. Are there any exceptions?

1. Approximately how much NaNO3 will dissolve in 100 g of water at 300C?

1. Approximately how much HCl will dissolve in 100 g of water at 200C?

1. How much NaCl will saturate 100 g of water at 500C?

1. Which of the solids in the graph does temperature seem to have the least effect on its solubility? Explain how you know.

1. How much does the solubility of NH3 decrease between 100C and 900C?

1. Which substance is least soluble at 0oC?

1. Thirty grams of KCl are dissolved in 100 mL of water at 45oC. How many additional grams of KCl are needed to make the solution saturated at 80oC?

1. If 50. mL of water that is saturated with KClO3 at 25oC is slowly evaporated to dryness, how many grams of the dry salt would be recovered?

1. Determine whether each of the following is saturated, unsaturated or supersaturated:

9. 95 g of NaNO3 in 100 g of water at 300C

9. 41 g NH3 in 100 g of water at 200C

9. 60 g of NH4Cl in 100 g of water at 500C

1. A saturated solution of potassium nitrate is prepared at 60oC using 100.mL of water. How many grams of solute will precipitate out of solution if the temperature is suddenly cooled to 30oC?

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 most and least 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

SO42−

2.65

7.6

Mg2+

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. 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?

8. Find [HCl] if 25.0mL concentrated HCl (12.0M) is diluted to a final volume of 2.00L.

9. How would you make 1.00L of a 0.50M NH3 solution given 15M NH3 solution?

10. If 25mL of a 0.2M Li3PO4 solution was poured into 130mL of pure water, calculate [Li+].

11. 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

12. 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) sulfateKsp = 1.30 × 10-8

b) silver carbonate Ksp = 6.2 × 10-12

c) magnesium hydroxideKsp = 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.

b)[SO42−] is reduced.

c)The solubility of BaSO4 is increased.

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 iodide ion and 8.4 × 10−4 M chloride ion. What concentration of added mercury (I) ion would cause precipitation of mercury (I) chloride and mercury (I) iodide? Which of these two mercury compounds would precipitate first? The Ksp for mercury (I) chloride is 1.4 x 10-18; the Ksp for mercury (I) iodide is 4.5 x 10-29.

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?Ksp = 1.1 × 10−12

17

8.310

-

´

[][]

AgI

+-

=