Unit 4
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Unit 4. Types of Chemical Reactions and Solution Stoichiometry Upon successful completion of this unit, the students should be able to: 4.1 Define the terms solute, solvent, and solution. 1.
Define solute.
4.2 Define the terms strong, weak, and non-electrolyte and solve related problems. 1.
Define strong electrolyte.
2.
Describe the difference between the terms slightly soluble and weak electrolyte.
3.
The compound AB, a weak electrolyte, is dissolved in water. Which picture below represents the best molecular level representation of AB in water (note: water molecules are present by not shown). a)
b) AB
AB
AB
B-
AB
AB
A+
AB
AB
AB
c)
d) +
A
A+
BB-
A+
B-
-
A+
AB
B-
4.
AB
AB
AB
B
+
A
A+ BB-
-
B +
A
A+ A+
B- A+
B-
A
Write the equation for the dissociation of the following salts in water: a. b. c. d.
LiNO3 CaSO4 Mg(C2H3O2)2 Na3PO4
4.3 Define Arrhenius acid and base. 1.
B-
+
Define Arrhenius acid and write a chemical equation demonstrating the definition.
4.4 Define and list common examples of: a. strong acids (HCl, HBr, HI, H2SO4, HNO3, HClO4) b. weak acids (H3PO4, HF, HC2H3O2 [CH3COOH]) c. strong bases (Group 1 hydroxides [LiOH, NaOH, KOH, RbOH, CsOH]; other hydroxides are strong bases in the sense that they completely dissociate into ions when dissolved in water, but most are typically not very soluble in water to begin with). d. weak bases (NH3). 1.
List two weak acids which ionize to produce a polyatomic anion.
2.
Define strong base.
4.5 List and/or recognize all ionic compounds and strong acids as examples of strong electrolytes. 1.
What is wrong with the following representation of sodium chloride dissolved in water.
4.6 List and/or recognize weak acids and weak bases as examples of weak electrolytes. 1.
Complete the table. Name
Formula
Acid or Base?
Strong or Weak?
sulfuric acid
________
___________
______________
____________
NH3
___________
______________
4.7 List and/or recognize the following common examples of nonelectrolytes: sugars (C6H12O6, C12H22O11), alcohols (CH3OH, CH3CH2OH), and water. 1.
Indicate if the following substances are strong electrolytes, weak electrolytes or nonelectrolytes. a. b. c. d. e. f.
NiSO4 HF C2H5OH HClO4 NH3 KOH
4.8 State the meaning of molarity. 1.
A bottle in the chemical storeroom is labeled 6.0 M HCl. Describe the type of information this label provides.
4.9 Calculate the molarity of a solution (or of the ions in solution of a strong electrolyte) and solve related problems. 1.
Which of the following solutions has the highest concentration of nitrate ions? (circle) 0.100 M Ca(NO3)2
0.150 M KNO3
0.075 M Mg(NO3)2
0.075 M Al(NO3)3
2.
Calculate the molarity of a solution that has 5.035 g FeCl3 dissolved in enough water to prepare 500. mL of solution.
3.
Determine the molarity of sulfate ions in 0.15 M Al2(SO4)3.
4.
Calculate the molarity of a solution made by dissolving 0.135 mol K2SO4 in enough water to form exactly 850. mL of solution.
4.10 Solve molarity problems involving calculating quantities of solute or solution. 1.
You want to prepare a 1.55 M solution of iron (III) bromide. How many grams of iron (III) bromide would you have to add to make 0.35 liters of this 1.55 M solution?
2.
How many moles of HNO3 are present in 25.0 mL of a 2.00 M HNO3 solution?
3.
How many mL of 0.350 M CuSO4 contain 2.00 g solute?
4.
An experiment calls for the addition of 0.170 g of sodium hydroxide in aqueous solution. How many milliliters of 0.150 M NaOH should be added?
4.11 Describe the proper techniques for solution preparation. 1.
What, if anything, is wrong with the following question: Calculate the milliliters of H2O which should be added to 17.5 g of NaCl to produce a 3.0 M NaCl solution?
2.
Describe how you would prepare 1.0 L of 0.25 M KMnO4 from a 3.0 M stock solution. Show all calculations AND describe your experimental procedures.
4.12 Solve problems involving dilutions of solutions. 1.
A 60.0 g sample of NaOH is dissolved in water, and the solution is diluted to give a final volume of 3.00 L. What is the molarity of the final solution?
2.
a.
b.
Your friend goes on vacation and asks you to fill her hummingbird feeder with a 1.50 M sucrose solution while she is gone. She gives you 525 mL of a stock solution, which contains 612 g of sucrose (C12H22O11). How much water do you need to add to the stock solution to dilute it to the concentration you need? Assume additive volumes. In order to arrive at an answer in part “a”, why was it necessary to assume additive volumes in order to determine how much water needed to be added? Explain.
4.13 Solve stoichiometry problems involving reactants and/or products in solution. 1.
How many milliliters of 0.62 M Ca(OH)2 are needed to produce 3.4 grams of H2O according to the following equation? Assume that HCl is in excess. 2 HCl
2.
+
Ca(OH)2 Æ
CaCl2
+
2H2O
How many mL of 0.136 M NaOH are required to react with the H2SO4 in 10 mL of a 0.202 M solution? 2NaOH + H2SO4 Æ Na2SO4 + 2H2O
3.
A chemical engineer determines the mass percent of iron in an ore sample by converting the Fe to Fe2+ in acid and then reacting the Fe2+ with MnO4-. A 1.1081 g sample was dissolved in acid and then reacted with 39.32 mL of 0.03190 M KMnO4. The balanced equation is: 8H+(aq) + 5Fe2+(aq) + MnO4-(aq) Æ 5Fe3+(aq) + Mn2+(aq) + 4H2O(l) Calculate the mass percent of iron (grams of iron per 100 gram of ore sample) in the ore.
4.
Alcohol levels in blood can be determined by reacting the alcohol with potassium dichromate according to the balanced equation: C2H5OH(aq) + 2Cr2O72-(aq) + 16H+(aq) Æ 2CO2(g) + 4Cr3+(aq) + 11H2O(l) alcohol What is the mass percent (grams of alcohol per 100 grams of sample) of the alcohol in the blood if 8.76 mL of 0.04988 M K2Cr2O7 is required to react with a 10.002 g sample of blood?
4.14 Recognize and describe the nature of a precipitation reaction, acid-base reaction, and redox reaction. 1.
Label the following reaction as precipitation, acid-base, or redox. 2Al(s)
2.
+
6HCl(aq) Æ
2AlCl3(aq)
+
3H2(g)
Classify the following reaction as precipitation, acid-base neutralization, or redox. S8(s) +
8O2(g) Æ
8SO2(g)
3.
Identify the following chemical reactions as either an acid-base reaction, a precipitation reaction or an oxidation-reduction reaction. (Then write your answer in the blank to the right of the chemical equation.) a. H2SO4 (aq) + 2LiOH (aq) Æ Li2SO4 (aq) + 2H2O (l)
Rxn: ___________________
b. Pb(NO3)2 (aq) + MgCl2 (aq) Æ PbCl2 (s) + Mg(NO3)2 (aq) Rxn: __________________ c. 2C3H6 (g) + 9O2 (g) Æ 6CO2 (g) + 6H2O (l)
Rxn: ___________________
d. NiBr2 (aq) + Ca(OH)2 (aq) Æ Ni(OH)2 (s) + CaBr2 (aq) Rxn: ____________________ 4.
5.
Classify the following reactions as precipitation (P), acid-base neutralization (AB), or redox (R). _______
S8(s) + 8O2(g) Æ 8SO2(g)
_______
2NaCl(aq) + Pb(NO3)2(aq) Æ 2NaNO3(aq) + PbCl2(s)
_______
CH4(g) + 2O2(g) Æ CO2(g) + 2H2O(l)
_______
HCN(aq) + LiOH(aq) Æ LiCN(aq) + H2O(l)
Differentiate between an acid-base reaction and a precipitation reaction. Can a reaction be both? If so, give an example.
4.15 Learn and be able to apply the following solubility rules for ionic compounds (and appreciate that these are general rules; even if two compounds are labeled soluble, the extent to which one of the compounds is soluble is likely different than the other). Soluble Ionic Compounds: 1. All common compounds of Group 1 ions (Li+, Na+, K+ etc…) and the ammonium ion (NH4 +) are soluble. 2. All common nitrates (NO3 -), acetates (CH3COO -), chlorates (ClO3-), and perchlorates (ClO4 -) are soluble. 3. Most common chlorides (Cl -), bromides (Br -), and iodides (I -) are soluble, except those of Ag +, Pb 2+, and Hg2 2+. 4. All common sulfates (SO4 2-) are soluble, except those of Ca 2+, Sr 2+, Ba 2+, Ag +, Hg2 2+, and Pb 2+. Insoluble Ionic Compounds: 1. All common metal hydroxides are insoluble, except those of Group 1 and the larger members of Group 2 (starting with Ca 2+). 2. All common carbonates (CO3 2-), phosphates (PO4 3-), and sulfides (S2-) are insoluble, except those of Group 1 and NH4 +.
1.
Would you expect the following compounds to be soluble in water? a.
SrSO4
b.
Ca(NO3)2
c.
(NH4)2CO3
d.
Cu(OH)2
2.
The following is an excerpt from an article in the Detroit Free Press (July, 27, 2004): “Two area lawmakers are calling for a statewide investigation into the use of lead pipes in public water systems. It's become an issue in Lansing, where the Board of Water and Light is voluntarily replacing lead pipes connected to as many as 14,000 properties over the next 10 years. Water is free of lead when it leaves the utility's plants but can become contaminated if it sits in lead pipes.” Imagine a similar dilemma occurred here in Champaign and the mayor of Champaign asks for your help to determine if the tap water in Champaign contains considerable amounts of Pb2+. Suggest a simple test the city could use to test for the presence of Pb2+?
3.
Will a precipitation reaction occur when aqueous solutions of FeCl2 and KOH are mixed? (answer yes or no)
4.
How would you prepare ZnCrO4 by a precipitation reaction? (in other words, list exactly the aqueous solutions of what two compounds would have to be mixed to yield ZnCrO4 as a precipitate) Assume you have an aqueous solution of an unknown salt. Treatment of the solution with dilute NaOH, Na2SO4, and KCl produces no precipitate. Which of the following cations might the unknown solution contain? (circle choices)
5.
Ag+ 6.
Cs+
Ba2+
NH4+
Criticize this statement. “All common chlorides (Cl -), bromides (Br -), and iodides (I -) are soluble, except those of Ag +, Pb 2+, Cu +, Hg2 2+. Therefore, carbon tetrachloride (CCl4) must be soluble in water”.
4.16 Predict products and write a balanced molecular equation and net ionic equation for precipitation reactions and acid-base reactions which are double displacement reactions. 1.
Decide whether a precipitate will form when the following solutions are mixed. If a precipitate forms, write a net ionic equation for the reaction. a. b.
barium nitrate and chromium (III) sulfate (NH4)2S and K2CO3
2.
Write a net ionic equation for the neutralization of HCN(aq) with NaOH(aq).
3.
Complete the following equation and then write a balanced net ionic equation. HBr(aq)
4.
+
Ca(OH)2 Æ
Give the chemical formulas for the products (be sure to indicate the phases for these products) that would be produced when the following reactants are mixed together. If no reaction would occur, write “No Reaction” on the line given for the products. a. aqueous potassium phosphate + Aqueous zinc chloride Æ ?? products (Write below) Products: ________________________________________________ b. aqueous hydrochloric acid + aqueous calcium hydroxide Æ ?? products (Write below) Products: ________________________________________________
c. aqueous ammonium bromide + aqueous lithium carbonate Æ ?? products (Write below) Products: ________________________________________________ d. aqueous strontium nitrate + aqueous aluminum sulfate Æ ?? products (Write below) Products: ________________________________________________ 5.
Write a balanced net ionic equation for the reaction between aqueous acetic acid and aqueous calcium hydroxide.
4.17 Identify spectator ions in a chemical reaction. 1.
Identify the spectator ions in problems 1, 2, 3 and 5 under the previous objective.
2.
Which of the choices below is NOT as spectator ion in the following reaction? BaCO3(s) + HCl(aq) Æ BaCl2(aq) + H2O(l) + CO2(g) a.
Ba2+
b.
CO32-
c.
H+
d.
all of these
4.18 Define titration and perform calculations related to titrations. 1.
Sodium carbonate, Na2CO3, is a good compound to use to standardize acid solutions. If 42.43 mL of HCl is required to completely react with 0.2844 g of Na2CO3, what is the molar concentration of the acid? Na2CO3
+ 2HCl Æ 2NaCl
+
CO2
+
H2O
2.
If 25 mL of 0.100 M HNO2 is titrated with KOH, how many moles of KOH are required to react with the HNO2.
3.
You titrate a 15.10 ml solution of sodium hydroxide with a 0.567 M solution of sulfuric acid. If 25.43 ml of sulfuric acid is needed to completely neutralize the sodium hydroxide, what is the molarity of the sodium hydroxide?
4.
Define titration.
4.19 State and apply the following rules for assigning oxidation numbers: 1. For free elements, the oxidation number is zero. Examples: Na (0), O2 (0), S8 (0). 2. For monatomic ions, the oxidation number is the charge on the ion. Examples: Na+ (+1), Cl- (-1), Al3+ (+3). 3. For Group 1A elements, the oxidation number is always +1 when they are in their ionic form. Examples: Na+ (+1), Rb+ (+1). 4. For Group 2A elements, the oxidation number is always +2 when they are in their ionic form. Examples: Ca2+ (+2), Ba2+ (+2).
5. For oxygen, the oxidation number is usually a -2. Exceptions: as a free element it is zero; for peroxides it’s a -1 (examples of peroxides: Na2O2, H2O2); for superoxides it’s -1/2 (examples of superoxides: RbO2, CsO2); with fluorine it’s positive (examples of oxyfluorine compounds: OF2, O2F2). 6. For hydrogen, the oxidation number is +1 when combined with other nonmetals and -1 when combined with metals. 7. For fluorine, the oxidation number is always a -1 (except when it is a free element). 8. When two elements are covalently bonded (usually a bond between two nonmetals), the element with the greater electronegativity is assigned a negative oxidation number equal to its charge in simple ionic compounds of that element. Example: In PCl3, Cl is more electronegative than P and since Cl forms a -1 charge in simple ionic compounds (such as NaCl, MgCl2), Cl is assigned an oxidation number of -1 in PCl3. 9. The sum of the oxidation numbers equals the charge on the molecule or ion. 1.
Give the oxidation number of the indicated element: a.
2.
O in OF2
b.
S in SO42-
Assign oxidation numbers to the indicated element in the following species: a. c.
N in NO2Cr in Al2(Cr2O7)3
b. d.
H in KH O in K2O2
4.20 Define oxidation, reduction, reducing agent, and oxidizing agent and solve related problems. 1.
In a redox reaction, a substance acts as an oxidizing agent. Has it gained or lost electrons?
2.
Where in the Periodic Table would the best oxidizing agents be found?
3.
Define oxidation.
4.21 Identify the following in an oxidation-reduction reaction: element getting oxidized, element getting reduced, oxidizing agent, reducing agent. 1.
Consider the following reaction: Cr(OH)3(s) + a. b.
2.
ClO3-(aq) Æ CrO42-(aq) + Cl-(aq)
What element is being oxidized? What is the oxidizing agent?
Consider the following reaction: S2O32-(aq) What is the reducing agent?
+
I2(aq) Æ
S4O62-(aq)
+
I-(aq)
3.
For the chemical reaction given below, determine what is oxidized, reduced, the oxidizing agent and the reducing agent and then write your answers on the appropriate line below the equation. 8 H+(aq) + 3 CH3OH(aq) + Cr2O7 2-(aq) Æ 3 CH2O(aq) + 2 Cr 3+(aq) + 7 H2O (l)
Oxidized: ________________________
Oxidizing Agent: ________________________
Reduced: ________________________
Reducing Agent: _________________________
Additional Unit 4 Sample Questions: 1.
Aqueous hydrochloric acid reacts with aqueous barium hydroxide to form aqueous barium chloride and water. If 1.25 Liters of 1.35 M hydrochloric acid reacts with 0.75 Liters of 1.25 M barium hydroxide, how many moles of water could be produced?
2.
When 30.0 mL of 0.657 M silver nitrate is mixed with excess calcium chloride, what is the mass of silver chloride solid that forms?
3.
Which of the following aqueous solutions is the best electrical conductor? (circle) 0.10 M CH3CH2OH
0.10 M C6H12O6
0.10 M KBr
0.10 M HC2H3O2
4.
a)
How many grams of solid will form if 23.7 mL of 0.179 M sodium sulfate is mixed with 30.6 mL of 0.130 M lead (II) nitrate? b) Determine the concentration of each ion remaining in the solution after the reaction has occurred.
5.
You have exposed electrodes of a light bulb to a reasonably concentrated solution of H2SO4 (just like the conductivity experiment in Figure 4.4 on page 137 of the textbook). a. b.
c.
Will the light bulb be off, dim, or bright? You now slowly add a solution of Ba(OH)2 to the H2SO4 solution and eventually the light bulb grows dim and finally goes off. Explain why the light bulb went off by using a net ionic equation and a written explanation. You now add even more Ba(OH)2 solution and the light comes back on again. Explain why.
Define solute.
4.2 Define the terms strong, weak, and non-electrolyte and solve related problems. 1.
Define strong electrolyte.
2.
Describe the difference between the terms slightly soluble and weak electrolyte.
3.
The compound AB, a weak electrolyte, is dissolved in water. Which picture below represents the best molecular level representation of AB in water (note: water molecules are present by not shown). a)
b) AB
AB
AB
B-
AB
AB
A+
AB
AB
AB
c)
d) +
A
A+
BB-
A+
B-
-
A+
AB
B-
4.
AB
AB
AB
B
+
A
A+ BB-
-
B +
A
A+ A+
B- A+
B-
A
Write the equation for the dissociation of the following salts in water: a. b. c. d.
LiNO3 CaSO4 Mg(C2H3O2)2 Na3PO4
4.3 Define Arrhenius acid and base. 1.
B-
+
Define Arrhenius acid and write a chemical equation demonstrating the definition.
4.4 Define and list common examples of: a. strong acids (HCl, HBr, HI, H2SO4, HNO3, HClO4) b. weak acids (H3PO4, HF, HC2H3O2 [CH3COOH]) c. strong bases (Group 1 hydroxides [LiOH, NaOH, KOH, RbOH, CsOH]; other hydroxides are strong bases in the sense that they completely dissociate into ions when dissolved in water, but most are typically not very soluble in water to begin with). d. weak bases (NH3). 1.
List two weak acids which ionize to produce a polyatomic anion.
2.
Define strong base.
4.5 List and/or recognize all ionic compounds and strong acids as examples of strong electrolytes. 1.
What is wrong with the following representation of sodium chloride dissolved in water.
4.6 List and/or recognize weak acids and weak bases as examples of weak electrolytes. 1.
Complete the table. Name
Formula
Acid or Base?
Strong or Weak?
sulfuric acid
________
___________
______________
____________
NH3
___________
______________
4.7 List and/or recognize the following common examples of nonelectrolytes: sugars (C6H12O6, C12H22O11), alcohols (CH3OH, CH3CH2OH), and water. 1.
Indicate if the following substances are strong electrolytes, weak electrolytes or nonelectrolytes. a. b. c. d. e. f.
NiSO4 HF C2H5OH HClO4 NH3 KOH
4.8 State the meaning of molarity. 1.
A bottle in the chemical storeroom is labeled 6.0 M HCl. Describe the type of information this label provides.
4.9 Calculate the molarity of a solution (or of the ions in solution of a strong electrolyte) and solve related problems. 1.
Which of the following solutions has the highest concentration of nitrate ions? (circle) 0.100 M Ca(NO3)2
0.150 M KNO3
0.075 M Mg(NO3)2
0.075 M Al(NO3)3
2.
Calculate the molarity of a solution that has 5.035 g FeCl3 dissolved in enough water to prepare 500. mL of solution.
3.
Determine the molarity of sulfate ions in 0.15 M Al2(SO4)3.
4.
Calculate the molarity of a solution made by dissolving 0.135 mol K2SO4 in enough water to form exactly 850. mL of solution.
4.10 Solve molarity problems involving calculating quantities of solute or solution. 1.
You want to prepare a 1.55 M solution of iron (III) bromide. How many grams of iron (III) bromide would you have to add to make 0.35 liters of this 1.55 M solution?
2.
How many moles of HNO3 are present in 25.0 mL of a 2.00 M HNO3 solution?
3.
How many mL of 0.350 M CuSO4 contain 2.00 g solute?
4.
An experiment calls for the addition of 0.170 g of sodium hydroxide in aqueous solution. How many milliliters of 0.150 M NaOH should be added?
4.11 Describe the proper techniques for solution preparation. 1.
What, if anything, is wrong with the following question: Calculate the milliliters of H2O which should be added to 17.5 g of NaCl to produce a 3.0 M NaCl solution?
2.
Describe how you would prepare 1.0 L of 0.25 M KMnO4 from a 3.0 M stock solution. Show all calculations AND describe your experimental procedures.
4.12 Solve problems involving dilutions of solutions. 1.
A 60.0 g sample of NaOH is dissolved in water, and the solution is diluted to give a final volume of 3.00 L. What is the molarity of the final solution?
2.
a.
b.
Your friend goes on vacation and asks you to fill her hummingbird feeder with a 1.50 M sucrose solution while she is gone. She gives you 525 mL of a stock solution, which contains 612 g of sucrose (C12H22O11). How much water do you need to add to the stock solution to dilute it to the concentration you need? Assume additive volumes. In order to arrive at an answer in part “a”, why was it necessary to assume additive volumes in order to determine how much water needed to be added? Explain.
4.13 Solve stoichiometry problems involving reactants and/or products in solution. 1.
How many milliliters of 0.62 M Ca(OH)2 are needed to produce 3.4 grams of H2O according to the following equation? Assume that HCl is in excess. 2 HCl
2.
+
Ca(OH)2 Æ
CaCl2
+
2H2O
How many mL of 0.136 M NaOH are required to react with the H2SO4 in 10 mL of a 0.202 M solution? 2NaOH + H2SO4 Æ Na2SO4 + 2H2O
3.
A chemical engineer determines the mass percent of iron in an ore sample by converting the Fe to Fe2+ in acid and then reacting the Fe2+ with MnO4-. A 1.1081 g sample was dissolved in acid and then reacted with 39.32 mL of 0.03190 M KMnO4. The balanced equation is: 8H+(aq) + 5Fe2+(aq) + MnO4-(aq) Æ 5Fe3+(aq) + Mn2+(aq) + 4H2O(l) Calculate the mass percent of iron (grams of iron per 100 gram of ore sample) in the ore.
4.
Alcohol levels in blood can be determined by reacting the alcohol with potassium dichromate according to the balanced equation: C2H5OH(aq) + 2Cr2O72-(aq) + 16H+(aq) Æ 2CO2(g) + 4Cr3+(aq) + 11H2O(l) alcohol What is the mass percent (grams of alcohol per 100 grams of sample) of the alcohol in the blood if 8.76 mL of 0.04988 M K2Cr2O7 is required to react with a 10.002 g sample of blood?
4.14 Recognize and describe the nature of a precipitation reaction, acid-base reaction, and redox reaction. 1.
Label the following reaction as precipitation, acid-base, or redox. 2Al(s)
2.
+
6HCl(aq) Æ
2AlCl3(aq)
+
3H2(g)
Classify the following reaction as precipitation, acid-base neutralization, or redox. S8(s) +
8O2(g) Æ
8SO2(g)
3.
Identify the following chemical reactions as either an acid-base reaction, a precipitation reaction or an oxidation-reduction reaction. (Then write your answer in the blank to the right of the chemical equation.) a. H2SO4 (aq) + 2LiOH (aq) Æ Li2SO4 (aq) + 2H2O (l)
Rxn: ___________________
b. Pb(NO3)2 (aq) + MgCl2 (aq) Æ PbCl2 (s) + Mg(NO3)2 (aq) Rxn: __________________ c. 2C3H6 (g) + 9O2 (g) Æ 6CO2 (g) + 6H2O (l)
Rxn: ___________________
d. NiBr2 (aq) + Ca(OH)2 (aq) Æ Ni(OH)2 (s) + CaBr2 (aq) Rxn: ____________________ 4.
5.
Classify the following reactions as precipitation (P), acid-base neutralization (AB), or redox (R). _______
S8(s) + 8O2(g) Æ 8SO2(g)
_______
2NaCl(aq) + Pb(NO3)2(aq) Æ 2NaNO3(aq) + PbCl2(s)
_______
CH4(g) + 2O2(g) Æ CO2(g) + 2H2O(l)
_______
HCN(aq) + LiOH(aq) Æ LiCN(aq) + H2O(l)
Differentiate between an acid-base reaction and a precipitation reaction. Can a reaction be both? If so, give an example.
4.15 Learn and be able to apply the following solubility rules for ionic compounds (and appreciate that these are general rules; even if two compounds are labeled soluble, the extent to which one of the compounds is soluble is likely different than the other). Soluble Ionic Compounds: 1. All common compounds of Group 1 ions (Li+, Na+, K+ etc…) and the ammonium ion (NH4 +) are soluble. 2. All common nitrates (NO3 -), acetates (CH3COO -), chlorates (ClO3-), and perchlorates (ClO4 -) are soluble. 3. Most common chlorides (Cl -), bromides (Br -), and iodides (I -) are soluble, except those of Ag +, Pb 2+, and Hg2 2+. 4. All common sulfates (SO4 2-) are soluble, except those of Ca 2+, Sr 2+, Ba 2+, Ag +, Hg2 2+, and Pb 2+. Insoluble Ionic Compounds: 1. All common metal hydroxides are insoluble, except those of Group 1 and the larger members of Group 2 (starting with Ca 2+). 2. All common carbonates (CO3 2-), phosphates (PO4 3-), and sulfides (S2-) are insoluble, except those of Group 1 and NH4 +.
1.
Would you expect the following compounds to be soluble in water? a.
SrSO4
b.
Ca(NO3)2
c.
(NH4)2CO3
d.
Cu(OH)2
2.
The following is an excerpt from an article in the Detroit Free Press (July, 27, 2004): “Two area lawmakers are calling for a statewide investigation into the use of lead pipes in public water systems. It's become an issue in Lansing, where the Board of Water and Light is voluntarily replacing lead pipes connected to as many as 14,000 properties over the next 10 years. Water is free of lead when it leaves the utility's plants but can become contaminated if it sits in lead pipes.” Imagine a similar dilemma occurred here in Champaign and the mayor of Champaign asks for your help to determine if the tap water in Champaign contains considerable amounts of Pb2+. Suggest a simple test the city could use to test for the presence of Pb2+?
3.
Will a precipitation reaction occur when aqueous solutions of FeCl2 and KOH are mixed? (answer yes or no)
4.
How would you prepare ZnCrO4 by a precipitation reaction? (in other words, list exactly the aqueous solutions of what two compounds would have to be mixed to yield ZnCrO4 as a precipitate) Assume you have an aqueous solution of an unknown salt. Treatment of the solution with dilute NaOH, Na2SO4, and KCl produces no precipitate. Which of the following cations might the unknown solution contain? (circle choices)
5.
Ag+ 6.
Cs+
Ba2+
NH4+
Criticize this statement. “All common chlorides (Cl -), bromides (Br -), and iodides (I -) are soluble, except those of Ag +, Pb 2+, Cu +, Hg2 2+. Therefore, carbon tetrachloride (CCl4) must be soluble in water”.
4.16 Predict products and write a balanced molecular equation and net ionic equation for precipitation reactions and acid-base reactions which are double displacement reactions. 1.
Decide whether a precipitate will form when the following solutions are mixed. If a precipitate forms, write a net ionic equation for the reaction. a. b.
barium nitrate and chromium (III) sulfate (NH4)2S and K2CO3
2.
Write a net ionic equation for the neutralization of HCN(aq) with NaOH(aq).
3.
Complete the following equation and then write a balanced net ionic equation. HBr(aq)
4.
+
Ca(OH)2 Æ
Give the chemical formulas for the products (be sure to indicate the phases for these products) that would be produced when the following reactants are mixed together. If no reaction would occur, write “No Reaction” on the line given for the products. a. aqueous potassium phosphate + Aqueous zinc chloride Æ ?? products (Write below) Products: ________________________________________________ b. aqueous hydrochloric acid + aqueous calcium hydroxide Æ ?? products (Write below) Products: ________________________________________________
c. aqueous ammonium bromide + aqueous lithium carbonate Æ ?? products (Write below) Products: ________________________________________________ d. aqueous strontium nitrate + aqueous aluminum sulfate Æ ?? products (Write below) Products: ________________________________________________ 5.
Write a balanced net ionic equation for the reaction between aqueous acetic acid and aqueous calcium hydroxide.
4.17 Identify spectator ions in a chemical reaction. 1.
Identify the spectator ions in problems 1, 2, 3 and 5 under the previous objective.
2.
Which of the choices below is NOT as spectator ion in the following reaction? BaCO3(s) + HCl(aq) Æ BaCl2(aq) + H2O(l) + CO2(g) a.
Ba2+
b.
CO32-
c.
H+
d.
all of these
4.18 Define titration and perform calculations related to titrations. 1.
Sodium carbonate, Na2CO3, is a good compound to use to standardize acid solutions. If 42.43 mL of HCl is required to completely react with 0.2844 g of Na2CO3, what is the molar concentration of the acid? Na2CO3
+ 2HCl Æ 2NaCl
+
CO2
+
H2O
2.
If 25 mL of 0.100 M HNO2 is titrated with KOH, how many moles of KOH are required to react with the HNO2.
3.
You titrate a 15.10 ml solution of sodium hydroxide with a 0.567 M solution of sulfuric acid. If 25.43 ml of sulfuric acid is needed to completely neutralize the sodium hydroxide, what is the molarity of the sodium hydroxide?
4.
Define titration.
4.19 State and apply the following rules for assigning oxidation numbers: 1. For free elements, the oxidation number is zero. Examples: Na (0), O2 (0), S8 (0). 2. For monatomic ions, the oxidation number is the charge on the ion. Examples: Na+ (+1), Cl- (-1), Al3+ (+3). 3. For Group 1A elements, the oxidation number is always +1 when they are in their ionic form. Examples: Na+ (+1), Rb+ (+1). 4. For Group 2A elements, the oxidation number is always +2 when they are in their ionic form. Examples: Ca2+ (+2), Ba2+ (+2).
5. For oxygen, the oxidation number is usually a -2. Exceptions: as a free element it is zero; for peroxides it’s a -1 (examples of peroxides: Na2O2, H2O2); for superoxides it’s -1/2 (examples of superoxides: RbO2, CsO2); with fluorine it’s positive (examples of oxyfluorine compounds: OF2, O2F2). 6. For hydrogen, the oxidation number is +1 when combined with other nonmetals and -1 when combined with metals. 7. For fluorine, the oxidation number is always a -1 (except when it is a free element). 8. When two elements are covalently bonded (usually a bond between two nonmetals), the element with the greater electronegativity is assigned a negative oxidation number equal to its charge in simple ionic compounds of that element. Example: In PCl3, Cl is more electronegative than P and since Cl forms a -1 charge in simple ionic compounds (such as NaCl, MgCl2), Cl is assigned an oxidation number of -1 in PCl3. 9. The sum of the oxidation numbers equals the charge on the molecule or ion. 1.
Give the oxidation number of the indicated element: a.
2.
O in OF2
b.
S in SO42-
Assign oxidation numbers to the indicated element in the following species: a. c.
N in NO2Cr in Al2(Cr2O7)3
b. d.
H in KH O in K2O2
4.20 Define oxidation, reduction, reducing agent, and oxidizing agent and solve related problems. 1.
In a redox reaction, a substance acts as an oxidizing agent. Has it gained or lost electrons?
2.
Where in the Periodic Table would the best oxidizing agents be found?
3.
Define oxidation.
4.21 Identify the following in an oxidation-reduction reaction: element getting oxidized, element getting reduced, oxidizing agent, reducing agent. 1.
Consider the following reaction: Cr(OH)3(s) + a. b.
2.
ClO3-(aq) Æ CrO42-(aq) + Cl-(aq)
What element is being oxidized? What is the oxidizing agent?
Consider the following reaction: S2O32-(aq) What is the reducing agent?
+
I2(aq) Æ
S4O62-(aq)
+
I-(aq)
3.
For the chemical reaction given below, determine what is oxidized, reduced, the oxidizing agent and the reducing agent and then write your answers on the appropriate line below the equation. 8 H+(aq) + 3 CH3OH(aq) + Cr2O7 2-(aq) Æ 3 CH2O(aq) + 2 Cr 3+(aq) + 7 H2O (l)
Oxidized: ________________________
Oxidizing Agent: ________________________
Reduced: ________________________
Reducing Agent: _________________________
Additional Unit 4 Sample Questions: 1.
Aqueous hydrochloric acid reacts with aqueous barium hydroxide to form aqueous barium chloride and water. If 1.25 Liters of 1.35 M hydrochloric acid reacts with 0.75 Liters of 1.25 M barium hydroxide, how many moles of water could be produced?
2.
When 30.0 mL of 0.657 M silver nitrate is mixed with excess calcium chloride, what is the mass of silver chloride solid that forms?
3.
Which of the following aqueous solutions is the best electrical conductor? (circle) 0.10 M CH3CH2OH
0.10 M C6H12O6
0.10 M KBr
0.10 M HC2H3O2
4.
a)
How many grams of solid will form if 23.7 mL of 0.179 M sodium sulfate is mixed with 30.6 mL of 0.130 M lead (II) nitrate? b) Determine the concentration of each ion remaining in the solution after the reaction has occurred.
5.
You have exposed electrodes of a light bulb to a reasonably concentrated solution of H2SO4 (just like the conductivity experiment in Figure 4.4 on page 137 of the textbook). a. b.
c.
Will the light bulb be off, dim, or bright? You now slowly add a solution of Ba(OH)2 to the H2SO4 solution and eventually the light bulb grows dim and finally goes off. Explain why the light bulb went off by using a net ionic equation and a written explanation. You now add even more Ba(OH)2 solution and the light comes back on again. Explain why.
