Unit 12 1  

HONORS CHEMISTRY

HARVARD-WESTLAKE

UNIT 12

Electrochemistry

Better living through chemistry:

C= Carbon Ho= Holium Co= Cobolt

La= Lanthunum Te= Tellurium

CHoCoLaTe

Unit 12 2  

Unit 12 3  

Galvanic Cell Practice Problems Draw a galvanic cell, utilizing a salt bridge, which consists of Cu in Cu2+ solution and Ag in Ag+ solution. Give the net cell reaction, E°, label the cathode, anode, and the direction of electron and ion flow.

Draw a galvanic cell, utilizing a salt bridge, which consists of Cu in Cu2+ solution and H2 in H+ solution. Give the net cell reaction, E°, label the cathode, anode, and the direction of electron and ion flow.

Unit 12 4  

Draw a galvanic cell, utilizing a salt bridge, involving the following cell reaction: Zn | Zn2+ || Co2+ | Co Give the net cell reaction, E°, label the cathode, anode, and the direction of electron and ion flow.

Draw a galvanic cell, utilizing a salt bridge, which consists of Cu in Cu2+ solution and Al in Al3+ solution. Give the net cell reaction, E°, label the cathode, anode, and the direction of electron and ion flow.

Unit 12 5  

Faraday’s Law In every electrochemical process, whether spontaneous or not, a certain amount of electric charge is transferred during the oxidation and reduction. The half-reactions we have written for electrode processes include the electrons which carry that charge. It is possible to measure the rate at which the charge is transferred with a device called an ammeter. An ammeter measures the current flowing through a circuit. The units of current are amperes (A) (amps, for short). Unlike a voltmeter, ammeters allow electrons to pass and essentially "clock" them as they go by. The amount of electric charge which has passed through the circuit can then be calculated by a simple relationship:

Charge = current x time OR Coulombs = amps x seconds This enables us to connect reaction stoichiometry to electrical measurements. The principles underlying these relationships were worked out in the first half of the 19th century by the English scientist, Michael Faraday. The key in applying this information to electrochemical systems is the Faraday constant: 96,500 C/mol e- (F). This proportionality connects the amount of charge to a definite number of electrons. Since each half-reaction gives the ratio of electrons to a particular substance, stoichiometric relationships can then be worked out from current and time measurements. In fact, some of the early determinations of Avogadro's number were done in this way. Once reasonably accurate relative atomic masses were known a simple plating cell could be used to find the relationship between current and time (therefore charge--or the number of electrons) and the amount of metal plated or dissolved.

Unit 12 6  

1. The electrolysis of molten MgCl2 is the last step in the isolation of magnesium metal from seawater. How many amps are needed to produce 35.6 g of Mg in 2.50 hours?

Unit 12 7  

2. A galvanic cell consisting of standard concentrations: Cu | Cu2+ || Ag+ | Ag, is connected to a small light bulb for 40.0 minutes. During that time the average current drawn by the bulb is 0.12 A. How many grams of copper dissolve from the anode during the time?

3. How many minutes does it take to produce 10.0 L of oxygen gas (measured at 755 mm Hg and 26.0 °C) by electrolyzing neutral water with a current of 1.3 A? What mass of hydrogen forms at the same time?

Unit 12 8  

4. A lead storage battery like one used to start your car contains lead and lead (IV) oxide electrodes in a sulfuric acid electrolyte. The electrode reaction are:

Anode: Pb + SO4

2- PbSO4 + 2 e- Cathode: PbO2 + 4 H+ + SO4

2- + 2 e- PbSO4 + 2 H2O

A typical automobile battery might be rated at “100 ampere-hours”. This means it can deliver a current of 1.0 A for 100 hours.

How many grams of Pb are oxidized for the full rating?

5. The violet vanadium (II) ion, V2+, can be produced by electrolysis of a solution containing the green V3+ ion. How long will it take to accomplish this in 0.35 L of a 0.50 M V2+ solution with a current of 0.26 A?

Unit 12 9  

Name:___________________________ Per.:____ Date:_______________

LAB: Galvanic Cells Purpose: to determine the ease of oxidation/reduction for a series of metals by measuring potentials in galvanic cells employing the metals; to examine the effect of changing concentration on cell potential. Method: I will make galvanic cell combinations of the various metals by placing the metals on a piece of filter paper which is soaked with the ionic solution of that metal. I will use NaNO3 solution to connect the wet spots together and then touch wires from the voltage probe to pairs of metals in order to measure the potential generated by each combination. I will also note which metal in each pair is attached to the “negative” [black] wire of the probe. In addition to the metal pairs using 1 M ionic solutions I will also measure the potential for Ag and Cu using 1 M and 0.0001 M concentrations of both ionic solutions to investigate the effect of changing the concentrations of the electrolytes. With the data from the experiment I will arrange the metals from most easily oxidized (most often “negative” in each combination) to least easily oxidized and compare this list to the Standard Reduction Potential Table. The data from the concentration experiments will allow me to decide how the spontaneous direction in a reaction relates to the measured potential and how changing the ionic concentrations affects cell potential. Data:

Unit 12 10  

Metal Pair Potential (V) “Negative” [Black]

Cu-Mg

Cu-Zn

Cu-Pb

Cu-Ag

Ag-Pb

Ag-Zn

Ag-Mg

Pb-Zn

Pb-Mg

Zn-Mg

Ag (0.0001)

Cu

Ag

Cu (0.0001)

Unit 12 11  

LAB: Galvanic Cells One of the characteristics of redox reactions is the transfer of electrons (similar to the exchange of protons in acid/base reactions). If the electrons can be directed through a piece of metal during this transfer, then an electric current is created. We generally call such a set-up a cell. A group of cells hooked together is a battery. Each substance has a different tendency to exchange electrons with another substance. To a large extent this tendency is indicated by the position of a substance on the activity series. The actual phenomenon is more complex than that, but for our purposes here, the object is to place the metals in this experiment into a series from most easily oxidized to most difficult to oxidize. A classic set-up for a galvanic or voltaic cell is shown below:

This looks very nice but in practice gives rather poor results and uses a lot of solutions. You can achieve far better results and use up less material by placing two pieces of metal at opposite ends of a strip of filter paper that has been wet with the appropriate solutions:

[the NaNO3 in the middle of the strip acts like the "salt bridge" or U-tube in the traditional set-up, allowing ions to migrate slowly across the paper, maintaining the electrical connection] If the wires from a voltmeter are touched to the metal pieces, the resulting voltage reading represents the potential of electrons to flow from one metal to the other. Further, the polarity of the metals (which one is connected to the + or - terminals of the voltmeter) indicates their relative positions in a series from easily oxidized [-] (the top of the activity series) to difficult to oxidize [+] (the bottom of the series). Micro-cell design adapted from: Establishing a Table of Reduction Potentials: Micro-voltaic Cells, Dan Holmquist, Jack Randall, Donald Volz, Chemistry with CBL, Vernier Software, 1995

Unit 12 12  

In addition to determining the relative strengths of the metals as oxidizing agents, you can also investigate the effect on the voltage of changing the solution concentrations for the cell combination of Cu with Ag. To make all of this easier (and quicker) you will use a piece of filter paper which has been cut as shown below:

By placing metal pieces on drops of their ionic solutions and “bridging” the gap in the center with NaNO3 you can easily set up a number of different reactions to investigate.

Preparing to experiment You will be provided with the following materials:

1. pre-cut filter paper as shown on the previous page 2. pieces of the following metals: Cu, Zn, Ag, Pb, Mg 3. 1 M solutions of Cu2+, Zn2+, Ag+, Pb2+, Mg2+ 4. 0.0001 M solutions of Cu2+ and Ag+ 5. 1 M solution of NaNO3

Design an experiment to measure the voltage and determine the polarity in cell produced from all possible combinations of the metal supplied. [compare the 0.0001 M Cu2+ only with the 1 M Ag+ and the 0.0001 M Ag+ only with the 1 M Cu2+] Technique 1. Measuring voltage For this experiment you can set up to read voltage continuously. Note that the RED connector on the voltage cable is + and the BLACK is -. A peculiar property of the voltage interface is that it indicates a small voltage even when nothing is attached to the wires!! Just ignore that. When you touch the leads to an actual source of voltage, you will get the correct reading if the negative (BLACK) connector is touching the metal at which oxidation occurs. Otherwise, you will get a very small (+ or -) voltage. Thus it is important to try reversing the leads if you get little response. None of the cells in this experiment should yield voltages in the mV range, so if yours do, try adding more solution around the metal or remoistening the "bridge" with NaNO3--or both.  

Unit 12 13  

2. Preparation of metal samples for best results Use the synthetic steel wool (DRY) to clean both sides of each metal piece. Note that you have 2 pieces of Cu and 2 pieces of silver. Prepare all of the metal pieces BEFORE you apply the solutions or the spots might be dry by the time you are ready. 3. Cleanup The paper can be discarded and the plastic sheet should be rinsed and dried. Since the Ag+ solution will tend to spread out over the paper, we strongly recommend that you handle the used pieces of metal and the paper with tweezers to avoid stains from the silver solution. The metal pieces should be rinsed with a squirt of distilled water and dried. Analysis 1. Make a table listing your metal combinations, the correct polarity for each metal in that combination and the

recorded voltage [this may already be part of your data, in which case you should skip this and go on to #2]. 2. Examine your results for the 1 M solutions and arrange the metals in a list from most easily oxidized (-) to most

easily reduced (+). Compare this list with the activity series in your book. Do they agree? 3. What is the effect on the voltage of changing the concentrations to 0.0001 for half of the cell? Do your results

suggest the direction in which the Cu/Ag reaction is spontaneous? (see below)

Cu(s) + 2 Ag+(aq) Cu2+(aq) + 2 Ag(s) 4. Refer to the table of standard reduction potentials in your text book (p. 629). Locate the Cu2+ Cu half reaction

(+0.34 v). Also note the Ag+ Ag half reaction at +0.80 v [note, these are the voltages for 1 M concentrations only]. How would you combine these two voltages to get approximately the voltage you measured for this combination? After you figure that out, determine the rest of the "expected" voltages that you measured in the lab. How do your results compare?

5. Each of these combinations resulted in a spontaneous reaction (positive voltage). The metal that was oxidized

(higher in the list you made earlier) is losing electrons and is called the anode. The other metal gains the electrons (becomes reduced) and is called the cathode. Note the relative positions of anodes and cathodes in the table on p. 615.

Finally, what do the actual reactions look like? You've seen one possible example above. Write the reactions that occurred in the correct direction for each cell combination.  

Unit 12 14  

Unit 12 15  

Name:___________________________ Per.:____ Date:_______________

LAB: Corrosion of Iron Purpose: to investigate the process of corrosion on iron and the effect of ion concentration and pH on the rate; to investigate the ability of some metals to “protect” iron from corrosion Method: I will place drops of various solutions on a clean strip of iron. The solutions contain “ferroxyl” indicator which is sensitive to Fe2+ (turns blue) as well as hydroxide ions (turns pink). The solutions are: distilled and tap water, NaCl, NaOH and HCl. I will make observations on the relative rate of blue color formation as well as specific areas of the drops which turn blue or pink. For the second part of the experiment I will attach clean strips of Cu, Sn, Zn and Mg to a clean strip of iron using alligator clips. I will then use the NaCl solution (with indicator) and place drops which overlap each metal strip and the iron strip. I will then record observations of blue color formation (indicating corrosion) and any other observations. I can use the data from the first part of the experiment to compare the effect of ion concentration in general by looking at the rates of corrosion with distilled and tap water as well as NaCl. The NaOH and HCl data will give me some idea of the effect of pH. Data from the second part of the experiment should allow me to decide what kinds of metals can “protect” iron from corrosion and why this might be so. Data: Table 1: Relative corrosion rate

Solution Relative Rate Observations

Distilled water

Tap water

0.1 M NaCl

0.1 M NaOH

0.1 M HCl

Unit 12 16  

Table 2: Corrosion inhibition

Metal Observations

Mg

Zn

Sn

Cu

Unit 12 17  

LAB: Corrosion of Iron So far we have looked at electrochemical processes which are useful in some way. Spontaneous reactions occurring in galvanic cells generate electric potential. Non-spontaneous reactions which we can force to occur in plating or other electrolytic cells also have their utility. But by far the most significant and undesirable spontaneous electrochemical process is corrosion. The most common example is the rusting of iron metal or iron alloys such as ordinary carbon steel. Corrosion costs this country billions of dollars each year in replacement and maintenance costs. From a purely economic point of view, then, it would be helpful to understand corrosion processes. Most metals react with oxygen, as a glance at the Standard Reduction Potential table will show. The reduction half-reaction of oxygen in moist air has a potential of +0.40 V:

O2 + 2 H2O + 4 e- 4 OH- With the notable exceptions of gold and silver, most common metals will form spontaneous redox couples with this half-reaction. In dry air corrosion is very slow or nonexistent (why?). One way to visualize what is happening in the corrosion of a metal like iron is to focus on what happens in and around a drop of water on the metal surface:

The oxygen in the air around the water slowly diffuses into the drop. It is the liquid water that brings the dissolved gas into contact with the metal. At the surface of the iron beneath the water drop the following half-reaction may occur:

Fe Fe2+ + 2 e-

The potential for this half-reaction is +0.44 v. Taken together the two processes constitute only the first step in the creation of what we commonly refer to as "rust". The net reaction for this first simple step is therefore:

2(Fe Fe2+ + 2 e-) O2 + 2 H2O + 4 e- 4 OH- 2 Fe + O2 + 2 H2O 2 Fe(OH)2 Iron(II) hydroxide is insoluble but its green color is almost never observed because it is ordinarily further oxidized by the oxygen:

2 Fe(OH)2 + ½ O2 + H2O ↔ 2 Fe(OH)3 The final product (when dry) has the reddish-brown flaky character we associate with rust. Although the reaction that produces Fe(OH)2 is technically an equilibrium process (all electrochemical processes are) the value of Kc is very large (>1099 at 298 K) and left unchecked it will go to completion. But the rate is relatively slow under normal atmospheric conditions and so it is still possible to manipulate the equilibrium somewhat by changing appropriate factors. The rates of corrosion reactions--and presumably their mechanisms--vary widely. Factors which influence the progress of the net reaction in the first step of the oxidation of iron may have an effect on the overall rate. The nature of the oxide product is also very important in affecting the extent of the corrosion. For example, aluminum is a very active metal, but its oxide, Al2O3, is very dense and forms a thin protective layer on the metal which discourages further corrosion. In contrast, iron rust (hydrated forms of Fe2O3 such as reddish-brown Fe(OH)3) is typically flaky and easily crumbles off to continually expose fresh metal for reaction.

Unit 12 18  

Although the mechanism for corrosion is not always well understood, it is clear how to prevent it. The surface of the metal must be protected from contact with oxygen. Paints, oils and other coatings are often used for this purpose. But it is also possible to take advantage of the electrochemical nature of the process by providing competition for the unwanted reaction in the form of an oxidation that requires less energy than the corrosion of the metal . For example, buried steel pipes can be protected from corrosion by attaching a piece of a more active metal to them. As the more active metal is oxidized, it continually supplies electrons to the steel pipe, thus preventing (or at least slowing) its oxidation [recall, oxidation is the loss of electrons]. What conditions favor corrosion? What kinds of metal combinations work best for corrosion protection? These are some of the questions you will investigate in this experiment. You can follow the progress (or lack of progress) of the reaction by using an interesting indicator called "ferroxyl". The solution contains phenolphthalein and potassium hexacyanoferrate(III)--K3[Fe(CN)6]. The hexacyanoferrate(III) ion reacts with Fe2+ ions to form an intense blue complex ion called prussian blue which has been used for years in carbon papers, blue-printing inks, typewriter ribbons, and artists pigments:

Fe2+ + Fe(CN)63- FeFe(CN)6

- The indicator is very sensitive and will detect extremely small quantities of Fe2+. The additional presence of phenolphthalein may help you to further follow the progress of the corrosion process. Preparing to experiment You will be provided with the following materials:

1. 0.10 M HCl solution* 2. the following solutions or liquids containing "ferroxyl":

distilled water tap water 0.10 M NaCl 0.10 M NaOH

3. "ferroxyl" indicator 4. strips of iron metal 5. thin strips of the following metals: Cu, Zn, Mg, Sn 6. four small alligator clips 7. steel wool 8. hand lens

*see Technique section Design a series of experiments to investigate the following:

the relative rate of corrosion with the available liquids or solutions the ability of various metals to "protect" iron from corrosion (see Technique section) (hint: pick the

solution from above that has the second fastest rate of reaction)

Unit 12 19  

Technique 1. Preparation of the iron strips It is very important to clean the iron strips thoroughly before beginning any part of the experiment. Use steel wool to clean the entire surface and be especially vigilant about any signs of rust. Wipe off the strip with a dry paper towel before experimenting to avoid interference by pieces of steel wool. It is equally important to rinse and dry the strips completely when you are finished with the experiment. 2. Using HCl with "ferroxyl" indicator The HCl solution does not have the indicator pre-mixed because of the remote possibility that small amounts of toxic HCN gas might be generated on long standing in the bottle. You should place a drop of HCl on the strip where you want it and then add a small drop of indicator to it. 3. Testing for corrosion protection by other metals In order to obtain meaningful results you must have good electrical contact between the thin metal strips and the iron strip. Cleaning the metals with steel wool is a very good idea. Then use the small alligator clips to attach the strips to the iron. Be sure the clips actually "bite" into the metal (especially with the Sn which is quite thin) and hold the thin strips against the iron. You can test all of the metals on one strip using an arrangement such as shown below:

Notice that the drops of whatever solution you choose should extend far enough to contact both the thin metal strip and the iron. The alligator clips have been placed on alternate sides to help keep the Fe strip level so that the drops will not run off. You will find that the thin metal strips are longer than the width of the Fe strip. You should fold the extra length over the edge to help hold the thin strip against the iron. 4. Clean up Once again, it is VERY important that the strip be rinsed thoroughly at the end of the experiment and dried completely. Iron will rust rapidly in the lab because of the chemical fumes present in small concentrations. If the thin metal strips have survived they can also be cleaned and dried for re-use. The alligator clips should be deposited in a beaker of dilute NH3(aq) in the fume hood for later cleaning and drying. The Chemicals Potassium hexacyanoferrate(III) is a solid consisting of ruby-red crystals. Aqueous solutions slowly decompose on standing and are somewhat light sensitive. It is used chiefly for blueprints, in photography, calico printing, in electroplating and for the tempering of iron and steel.

Unit 12 20  

Analysis

1. Which solution showed the slowest rate of corrosion? The fastest? How do the solutions in between rank? Explain

why you think the results turned out the way they did. Your explanation should take into account the electrochemical nature of corrosion and the net reaction for the first step of iron corrosion. Be sure to mention the sodium hydroxide solution in particular.

2. The blue color of the "ferroxyl" indicator tells you that Fe2+ has been produced. The pink color of phenolphthalein

indicates that the solution is becoming basic. Explain the appearance of the drops in which corrosion occurred. Specifically, why are the colors segregated the way they were?

3. Why is the sodium hydroxide solution pink? 4. Sodium chloride and calcium chloride are commonly used to de-ice roads in the eastern part of the country (by

lowering the freezing point of the slush). When this salty slush is splashed up onto automobiles it has an undesirable effect on them. Based on your observations, why is this so?

5. Explain your observations for the "protected" metal combinations in light of the activity series. 6. Magnesium is a very active metal which, like aluminum, rapidly develops an oxide coating that prevents further

corrosion AND masks the reaction of the metal with water (remember the Periodic Properties experiment?). When employed as a sacrificial metal, however, the preferential oxidation of magnesium provides a constant clean surface. What evidence did you see for the reaction of magnesium with water when used to protect the iron?

7. Go to the Standard Reduction Potential table in your text and find the water half-reaction that will combine with

the oxidation half-reaction of magnesium and account for all of the observations you made (including any color changes). Write the net reaction that occurs between magnesium and water.

[This protective aspect of magnesium versus iron actually has a relatively recent practical application that is not connected with corrosion. The U.S. Army developed (along with Zesto Therm company) a heat pack for warming the infamous MREs or "meal, ready-to-eat". The heater consists of magnesium metal, sodium chloride and powdered iron all mixed into a porous pad enclosed in cardboard. When water is poured onto this package and it is placed underneath an MRE it quickly heats up as the reaction you have observed takes place. ΔH = -355 kJ/mol Mg.]

Unit 12 21  

Name:___________________________ Per.:____ Date:_______________

LAB: Electrolysis of Aqueous Solutions Purpose: to use my observations from the electrolysis of a few aqueous solutions to attempt to formulate—and explain—a rule for predicting the products of this process Method: I will place each solution in a u-tube and add an appropriate indicator to the top of the solution column in each arm. For NaF I will use bromthymol blue on each side. CuBr2 requires no indicator. For CoCl2 I will use green dye in each arm. For the KI I will place both starch and phenolphthalein in each arm. I will record my observations of what happens at each electrode. With the observations I will attempt to identify potential products in each electrolysis and decide whether the ions from the compound have been oxidized/reduced or whether water has undergone these processes. Using the SRP table I will look for any pattern in my results that might enable me to establish a rule for predicting products. Data: Table 1: electrolysis observations

Solution Positive [Red] Negative [Black]

NaF

CuBr2

CoCl2

KI

Unit 12 22  

Unit 12 23  

LAB: Electrolysis of Aqueous Solutions You have seen that chemical reactions in solution can produce electricity. It might seem logical that electricity passed through a solution may result in a chemical reaction. In fact, many of you will have done this very thing before in another science class: the electrolysis of water to produce hydrogen and oxygen gas. For that process we might write:

electrical energy + H2O(ℓ) H2(g) + ½ O2(g) It can be shown that this is a redox reaction by inspecting the changes in oxidation number. It should therefore be possible to write half-reactions for the processes which occur at each electrode, cathode and anode. We use these terms to mean the same thing in an electrolytic cell as they meant in a galvanic cell. One possible way to investigate these changes is with an acid/base indicator. Since water contains small amounts of H+ and OH-, it may be that these ions are involved in the breakdown of water into its elements. In other words, the mechanism for the reaction written above might involve these ions. Reactions like this require some external source of electrons (we usually call that electricity!) in order to occur. A battery or power supply can act as the “electron pump". As you have noted in your experiments with galvanic cells (simple batteries) electrons emerge from the negative (-) electrode and return to the positive (+) electrode. Placing these electrodes into a solution makes the electrons available for redox reactions. You may also remember that in order to electrolyze water, some electrolyte dissolved in the water was necessary. This is true since there are very few free ions in pure water, only about 2 x 10-7 M total (why?). Does the nature of the electrolyte affect the products of the electrolysis? The questions posed above can be investigated in a series of simple experiments, but it would be wrong to surmise from this that electrode processes in such cells are simple. In fact just the opposite is true. A complete quantitative treatment of aqueous electrolytic cells is beyond the scope of this course. But the fundamentals--which you can investigate in the lab--are not. Preparing to experiment You will be provided with the following materials:

1. mini U-tube 2. power supply (set on 12 volts) 3. graphite electrodes 4. white plastic 5. 0.1 M NaF solution 6. 0.1 M CoCl2 solution 7. 0.1 M CuBr2 solution 8. 0.1 M KI solution 9. bromthymol blue indicator (use 1 drop in each arm) 10. phenolphthalein indicator (use 1 drop in each arm) 11. starch solution (use 1 drop in each arm) 12. food coloring dye (use 1 drop ONLY in each arm)

Design an experiment using bromthymol blue indicator to investigate the changes in a sodium fluoride solution when electricity passes through it. Design an experiment using food coloring dye to investigate the changes in CoCl2 solution when electricity passes through it.

Unit 12 24  

Design an experiment to investigate the changes in CuBr2 solution when electricity passes through it. Design an experiment using phenolphthalein indicator and starch solution to investigate the changes in a KI solution when electricity passes through it. In each case be sure to note down all visual observations such as changes in color, gas formation and relative rates of gas formation if gases form at both electrodes. Pre-lab take-home quiz These questions should be answered on a separate sheet of paper to be turned in on the day you do this experiment. 1. Show, by assigning oxidation numbers, that the decomposition of water into hydrogen and oxygen is a redox

reaction. 2. It is stated in the introduction that the total ionic concentration in pure water is about 2 x 10-7 M. Why is this? Technique 1. Using mini U-tubes as electrolysis cells The mini U-tubes are designed to keep the electrode reactions more or less separate but still allow ions to migrate through the solution and carry the electricity. You can fill a tube directly from the squeeze bottles and add the indicators with the droppers provided. Then attach the graphite electrodes to the alligator clips from the power supply. Insert the electrodes into the arms of the U-tube and make your observations. Be careful not to touch the electrodes together or you could damage the power supply. You also might get a nasty shock if you are careless. Be sure to note which electrode (+ or -) is in which arm when you record your observations. Not all the reactions happen equally fast. Be sure to give enough time to see any changes before you move on. The diagram below is provided for your reference. You will probably find it easier to support the U-tube in a 50 mL beaker while working. The white plastic provided can be used behind it in order to better observe changes.

2. Using starch to test for I2 Starch (ordinary clothes starch or even cornstarch would work just fine) provides a sensitive test for I2 in solution that is used often in the lab. It forms an intense blue complex molecule (sometimes appearing almost black or dark purple) in the presence of I2. Otherwise it is colorless.

Unit 12 25  

Analysis 1. Recall the colors of bromthymol blue in various solutions: acid yellow neutral green base blue

According to this, what ions must be produced at the + and - electrodes during the electrolysis of an aqueous solution of NaF (along with the hydrogen and oxygen gas)? Why does the indicator start out blue or bluish green in the NaF solution? (hint: what acid and base form NaF?)

2. Which gas (H2 or O2) was produced at which electrode (+ or -) in the sodium fluoride cell? Based on your visual

observations at each electrode, HOW CAN YOU TELL? 3. Find the appropriate half-reactions for this process in the Standard reduction potential table in your text (hint:

they are found at -0.83 v and +1.23 v). Combine them to give the overall reaction you observed (you will have to reverse one).

a. What is the net cell voltage? b. Which electrode reaction occurs at the anode? Which electrode reaction occurs at the cathode? c. Generalize a rule for the location of anodes and cathodes in the table for an electrolytic cell d. Why doesn't the sodium fluoride solution electrolyze spontaneously? e. Since sodium and fluoride ions do not appear in the reaction, what function do they serve?

4. In order for the reaction to occur, the net cell voltage must be positive. You added +12 v to the sodium fluoride

cell. Does that result in a net positive voltage? 5. Use your results for the electrolysis of CuBr2 solution to determine the following:

a. the half-cell reactions at each electrode b. the net cell reaction c. the cell voltage

[refer to the Standard reduction potential table in your text]

6. Which gas (H2 or O2) is produced in the KI cell? [hint: at which electrodes were these gases produced in earlier

cells?] 7. Considering your answer to #6, use your results for the electrolysis of KI solution to determine the following: a. the half-reactions at each electrode b. the net cell reaction c. the cell voltage

[refer to the Standard reduction potential table in your text]

Unit 12 26  

8. Why isn't potassium metal a product of this electrolysis? 9. Generalize a rule for "guessing" which of two competing electrode processes (for example, the reduction

of water to produce hydrogen gas or the reduction of Co2+ to produce cobalt metal) will occur in the electrolysis of an aqueous solution. [note: this "rule" will only be correct in dilute solutions such as the ones you used in this experiment; in concentrated solutions all bets are off]

10. Why do you think cobalt metal is produced instead of hydrogen gas in the electrolysis of CoCl2 solution? [hint:

compare the voltages in the table]. 11. What gas do you think was produced at the other electrode? HOW CAN YOU TELL? [hint: what does chlorine

do to dyes?]. How does this fit in with your generalization given in #9? Explain. 12. Use your results for the electrolysis of CoCl2 solution to determine the following:

a. the half-reactions at each electrode b. the net cell reaction c. the cell voltage

[refer to the Standard reduction potential table in your text]

Unit 12 27  

LAB: Electroplating One might question the truth of that quotation, but it does have something to do with the problem at hand. Consider the men who plundered the Aztec treasure (?). Gold-plated copper ornaments were a great disappointment to the Spaniards when melted down--and they are a surprise to us even today. Gold (along with silver) is one of the so-called "noble" metals. It is highly resistant to acid attack (only aqua regia is able to dissolve gold; this is a mixture of concentrated HCl and concentrated HNO3). In order to plate gold onto copper as the Aztecs apparently did, they would need a solution of gold. But presumably they had no concentrated acids, and gold solutions do not occur in nature. They had HCl (in their stomachs) but probably didn't use it. HNO3 was most likely unavailable. Silver presents a similar dilemma. It will dissolve in HNO3 (if you have some) and then the solution can be used to plate baser metals. But what if you don't have any nitric acid? Consider a solution of nitric acid. It contains three types of ions: H+, NO3

- and OH-. The [OH-] is so small we can neglect it here. Now consider a solution of more mundane materials: vinegar (which contains acetic acid) and saltpeter (saltpeter is either sodium or potassium nitrate, occurring in natural deposits). Neglecting the OH-, four ions are present in this solution: H+, NO3

-, CH3COO-, Na+. Notice that the building blocks of nitric acid are present, although the concentration of H+ would be small since acetic acid is a weak acid. But left alone long enough, the silver from native ore deposits would slowly dissolve. Heating would help. As for gold, adding ordinary table salt to the solution used to dissolve silver would provide the Cl- ion and give the mix of ions present in aqua regia. Of course, there is no way of knowing exactly how the Aztecs might have stumbled upon these processes--or even if they did. But it is important to point out that not all useful chemistry has to come off the stockroom shelves. As Pasteur said: chance favors the prepared mind. Electroplating is an important industrial process today. It is based on the concepts of redox chemistry in electrolytic and galvanic cells. Spontaneous plating is sometimes called deposition and produces coatings that are unsuitable for anything but arousing the curiosity of chemistry students. Plating in electrolytic cells can be controlled to produce tightly adhering coatings such as the chromium on car bumpers (well, old cars) or the silver on expensive flatware. Slower processes usually produce higher quality coatings, so the Aztecs--even though they undoubtedly used some sort of spontaneous process--may have gotten pretty good results simply because they had to wait since their solutions would have been so dilute! In electrolytic plating operations, the anode is made out of the metal to be plated and the cathode is the object on which the plating is to be done (why?). The plating solution may be an electrolyte containing ions of the metal to be plated, but it may contain other ions as well. An external voltage is applied and the quality of results depends not only on surface preparation but also on the rate which is controlled by the current and voltage. Practical plating is something of an art as well as a science and some apparently unrelated ingredients are often added because Someone once found by accident that they enhanced the quality of the final product. Quantitatively, spontaneous plating must be a stoichiometric process since we can write balanced equations to represent what happens. For example, a piece of copper wire immersed in silver nitrate will undergo the following reaction:

Cu(s) + 2 Ag+(aq) Cu2+(aq) + 2 Ag(s) Technically this is an equilibrium reaction, but in practice it goes nearly to completion (Kc is 3.9 x 1038 at 25 oC !!). Penny plating technique adapted from: The "Golden Penny" Demonstration, Steven H. Szczepankiewicz, Joseph F. Bieron, Mariusz Kozik, J. Chem Educ., 1995, vol. 72, No. 5, p. 386

Unit 12 28  

Electrolytic plating is somewhat different. While mass is conserved in the example above, charge is also conserved and this is a key concept for understanding what happens in an electrolytic plating cell. If the anode and cathode are made out of the same material, atoms of metal dissolve at one electrode and are plated onto the other. The number of atoms that do this is related to the amount of charge that passes through the cell. Electric charge is measured in units called coulombs. One coulomb (C) is equal to one ampere flowing for one second:

charge = current x time

Coulombs = amps x seconds Since the charge on an electron is known (1.602 x 10-19 C), it is possible to determine the relationship between charge and a mole of electrons:

1 mole of electrons 96,500 C This number is known as the Faraday (F) and it is a very useful proportionality. Consider the example below:

A 0.500 amp current flowing for exactly one hour is passed through a solution of nickel (II) ions.

a. How many grams of nickel metal will plate out? b. If this cell is hooked in series with another cell containing Ag+ ions [in series the same current passes through both cells], how many grams of Ag will plate at the same time?

First we need to know how much charge passed through the cell(s):

0.500 A x 3600 s = 1800 C

Now how many moles of electrons was that?

1800 C x -1 mol e

96500 C = 0.0187 mole e-

Since nickel (II) has a charge of +2, one atom of nickel will require 2 e- to plate:

0.0187 mole e- x -1 mole Ni2 mole e = 0.00935 mole Ni

This is 0.549 grams.

Since silver is plated by the same amount of charge, but each atom of silver requires only one electron to plate (Ag ions are +1), we should expect twice as many moles of silver atoms to plate:

0.0187 mole e- x -1 mole Ag1 mole e = 0.0187 mole Ag

Thus 2.02 grams of silver are produced. Note that the ratio based on charge affects the moles of silver and nickel. Their atomic masses influence the final mass plated.

Some authors (Chang, for one) call one mole of electrons a Faraday, but the meaning is still 96,500 C of charge.

Unit 12 29  

Plating experiments are rather boring and this one is no exception. There is not much to do as the process takes place (especially if your calculator is recording the data...) and there is generally not much to look at. However, there is time for a fascinating diversion. You know from the activity series that zinc, which is above copper, will be displaced by copper in solution. The voltage for this process is 1.10 v as you saw in the previous experiment. The opposite process, zinc displacing copper is non-spontaneous. Or is it? For many years chemistry teachers have done a demonstration in which they turn ordinary pennies into "gold". Traditionally the process has been done in a hot, very alkaline solution containing the ion Zn(OH)4

2- and some solid zinc. This combination of materials caused zinc metal to spontaneously plate on the copper. When the "silver" pennies were subsequently heated in a flame, the thin coating of zinc alloyed with the copper, producing yellow brass or "gold". Recently a pair of chemists investigated this process more closely. Their report appeared in the Journal of Chemical Education. They found that the process was very different from what had been assumed. In fact, the very alkaline solution was not required at all! Copper placed in a heated solution of Zn2+ and a small amount of metallic zinc will become plated with zinc metal!! How is this? According to their research the process is driven by the alloying which occurs at the very surface of the copper. Due to thermal agitation in a heated solution, some zinc ions manage to work their way into the spaces between the copper atoms and form a very thin layer of brass. The voltage for zinc metal plating on brass (as opposed to copper) is positive! Thus the "silver" appearance of the penny after heating is actually a type of brass (silver brass). When heated in a flame, the silver brass is converted to yellow brass ("gold"). You can do this yourself, using the plating solution and some granular zinc. Bring two of the shiniest (cleanest) pennies you can find. Place them in the beaker of solution provided on the hot plate and boil the solution. It takes about 10-15 minutes to plate a uniform layer of silver brass. Turning the pennies occasionally will help. When they are covered in silver brass, remove them and rinse in water. You can keep one "silver" penny and heat the other very carefully in a cool burner flame to change it to "gold". Too hot a flame will melt the brass coating, so be careful. This is just for fun, but it also points up the fact that not everything which we take for granted is happening they way we think. There are plenty of puzzles left to solve. Be sure to keep an eye on your experiment while you do this little bit of modern alchemy! Preparing to experiment You will be provided with the following materials:

1. two numbered zinc metal strips 2. steel wool 3. zinc plating solution (containing mainly ZnSO4) 4. a power supply (set to 1.5 volts) 5. a current probe 6. 100 mL beaker and electrode holder

In addition, beakers of distilled water and acetone will be available in the lab and the oven will be on. Design an experiment to compare the amount of electric charge flowing through a zinc plating cell with the mass loss and mass gain of the electrodes. Running the cell about 15-20 min should give sufficient mass changes to measure precisely on the balance. Measuring current each 30 seconds should be adequate.

Unit 12 30  

Technique 1. Measuring current The current probe consists of a small resistance shorted across the leads. Like a normal ammeter, it must be

hooked in series with the circuit. For your plating cell the circuit would look like this:

Note that the positive end of the power supply must be connected to the positive terminal of the ammeter (the red lead on the current probe). Electrons emerge from the negative terminal of the power supply and flow through the plating cell, through the ammeter, and back to the power supply, entering at the positive end.

Notice the level of the plating solution in the beaker. It should NOT touch the alligator clips that hold the

electrodes. Also you should not put the numbered ends of the electrodes into the solution or the numbers will eventually disappear!

2. Massing electrodes after plating The electrodes will be wet, of course, and you will want to mass them dry. The quickest way to do this is to gently

rinse them in distilled water by dipping, followed by a similar rinse in acetone (which has a very high vapor pressure at room temperature) followed by a short stay in the oven. Be careful not to scrape off any new material adhering to the electrodes.

Analysis 1. Calculate the mass changes for the two zinc electrodes in the plating cell.

a. How are they related? Does this seem correct? Why? b. Which electrode increased/decreased in mass relative to which terminal of the power supply (+ or -) it was

connected to? Why?

2. From your current data determine the average current that flowed through the plating cell for the duration of the experiment* and use this information to determine the amount of charge (in Coulombs) which passed through the cell.

3. From the number of Coulombs, calculate the expected mass change of zinc metal and compare this to your mass

change data. Comment on any difference, describing factors which you suspect may have contributed to error.

Unit 12 31  

Name:___________________________ Per.:____ Date:_______________

LAB: Electrochemical Production of Hydrogen Gas and Copper(II) ion Iodometric Titration of Copper(II) ion

Purpose: to determine the amount of copper produced by the electrolysis of an acidic solution using a copper anode and an inert cathode in a variety of ways and compare the results Method: After massing the copper electrode, I will begin the experiment by assembling the electrolysis apparatus as shown in the lab handout. I will measure the current every 15 seconds while collecting the hydrogen gas generated at the cathode (up to 25 mL). After recording the volume of gas collected (and the room temperature and pressure) I will drain the apparatus and transfer the liquid to a graduated cylinder, bringing the volume up to 100 mL with water. After drying the copper electrode I will remass it. 1.5 mL of the electrolysis solution will be combined with 1.5 mL of ammonia solution to produce a deeper blue solution for which I will measure the % transmittance. I will titrate the remaining 98.5 mL of solution with sodium thiosulfate solution, using starch as an indicator. The mass change of the electrode represents the amount of copper that dissolved (to one significant digit). I can determine the moles of hydrogen collected using the volume of gas, and the room temperature and pressure (n = PV/RT) and compare this to the moles of copper through the stoichiometry of the electrolysis reaction. From the time and current I can determine the moles of copper as well, since seconds x amps = coulombs and there are 96,500 coulombs per mole of electrons. Each copper atom oxidized loses two electrons. The absorbance (2-log %T) of the small sample treated with ammonia can be compared to the standard to determine the moles of copper that way. Finally, the titration data can also be used to determine the moles of copper through the stoichiometry of the reaction, using the volume and concentration of the sodium thiosulfate solution.

Unit 12 32  

Data: Table 1: Change in copper electrode mass

Mass, g

Before electrolysis

After electrolysis

Table 2: Hydrogen gas data

Room temperature, °C

Room pressure, mm Hg

Volume, mL

Table 3: % Transmittance data

% Transmittance

1.5 mL 0.010 M Cu2+ + 1.5 mL 6 M NH3

1.5 mL your solution + 1.5 mL 6 M NH3

Table 4: Titration data

Buret reading mL

Initial

Final

Table 5: Current data

Total Time, s

Average Current, A

Unit 12 33  

LAB: Electrochemical Production of Hydrogen Gas and Copper(II) ion Iodometric Titration of Copper(II) ion

Electrolysis of water solutions in which one electrode is made of a material more easily oxidized than water results in the dissolution of that electrode. This process is used industrially to purify metals, particularly copper. In that case, the copper which is oxidized is in the form of an impure anode. The resulting copper(II) ions then plate onto a pure Cu cathode, leaving impurities behind. In this experiment, the copper oxidized remains in the solution and the presence of acid allows H+ ion to be reduced at the relatively inert stainless steel cathode to form H2 gas. Thus in this electrolytic cell, an external voltage causes copper metal to be oxidized to copper(II) ion and H+ ion to be reduced to hydrogen gas. The amount of copper electrolyzed can be determined in a variety of ways. The electrode can be massed before and after. The electrical work involved can be used to determine the amount of Cu using Faraday's laws. The concentration of Cu2+ in solution can be determined with a colorimeter. Or the Cu(II) can be titrated with a suitable reducing agent. All of these methods are employed in this experiment:

(1) The mass change of the copper electrode is too small to permit precise determination on the centigram balances, but it can be used as a ballpark check on the overall results.

(2) The Faraday determination is similar to that done in the plating of zinc metal. (3) The copper(II) ion from the oxidation of the electrode colors the solution blue and could be determined

directly by a sensitive colorimeter, but the concentration is too small for our equipment. This is not an uncommon problem and there is a common solution: add something to the mixture which intensifies the color (or creates a new one). Aqueous ammonia reacts with copper(II) ion to form a deep blue complex ion, Cu(NH3)4

2+. A copper(II) solution with known concentration is provided so you can compare its absorbance when treated with NH3 to the absorbance of your sample when treated in the same way.

(4) To allow for comparison of the techniques and to illustrate the process of redox titration (in this case, iodometric titration), the copper(II) ion in solution will also be titrated as described on the facing page.

(5) Finally, the volume of hydrogen gas collected may also be used to confirm the copper calculations from Faraday's laws, using the ideal gas law.

Iodometric titration of transition metals is often accomplished using thiosulfate ion (S2O3

2-) as a reducing agent. In this case iodide ion (from solid KI) is used to reduce the copper(II) produced by the electrolysis to copper(I) which is insoluble with iodide. The iodide ion is oxidized to I2 which reacts with excess iodide in the solution to form the brown triiodide ion (I3

-) 2 Cu2+(aq) + 5 I-(aq) 2 CuI(s) + I3

-(aq) (1)

[the triiodide ion is considered equivalent to I- + I2] Titration of the triiodide ion (I3

-) is then begun by adding Na2S2O3 solution: I3

-(aq) + 2 S2O32-(aq) 3 I-(aq) + S4O6

2-(aq) (2) The brown color gradually fades to a light beige color as the titration proceeds. At this point, starch indicator is added, which causes the solution to turn blue-grey (I2 + starch blue-grey complex). Continued addition of thiosulfate ion results in the disappearance of the blue color as I3

- is reduced to I-. The final solution is white. Once the solid KI has been added to the solution the titration should be completed quickly (but not recklessly!!) for best results since copper(I) ion is readily oxidized in the air and this will tend to produce inaccurate results. Adapted from: Electrochemical Production of Hydrogen Gas and Cupric Ion. The Iodometric Titration of Copper, J. McCullough and H. Stone, Experimental General Chemistry, McGraw Hill, 1953

Unit 12 34  

Preparing to experiment You will be supplied with the following materials:

1. 6 M H2SO4 (use about 3 mL) 2. a copper electrode and a stainless steel electrode 3. a current probe 4. a colorimeter w/cuvette 5. a power supply (set on 12 volts) 6. 0.050 M Na2S2O3 solution 7. solid KI (use about 2.5 g) 8. 6 M NH3(aq) 9. 0.010 M standard copper(II) solution 10. starch solution (use about 3 mL) 11. a buret and clamp 12. magnetic stirrer/stirring bar 13. two 1 mL calibrated beral pipets 14. a 100 mL graduated cylinder

Design an experiment to electrolyze an acidic solution using a copper anode and a stainless steel cathode, collecting the hydrogen gas (about 25 mL). Determine [Cu2+] in the resulting solution by redox titration as described in the introduction, by the volume of H2 collected, the mass change of the copper electrode, and the absorbance of a sample of the copper(II) solution reacted with aqueous ammonia. Pre-lab take-home quiz These questions should be answered on a separate sheet of paper to be turned in on the day you do this experiment. 1. Write the balanced equation for the reaction that occurs during the electrolysis. 2. Add reactions (1) and (2) on page 1 to obtain the overall reaction of S2O3

2- with Cu2+ during the titration. 3. During a titration as described in the introduction, 15.5 mL of 0.15 M Na2S2O3 is used to reach the endpoint of a

copper solution. How many moles of copper are present? 4. In order for the copper electrode to dissolve during the electrolysis, which terminal of the power supply should it

be connected to (+ or -)? Technique 1. An electrolysis set-up:

 

 

 

  

Unit 12 35  

Notice that the electrode attached to the (+) terminal of the power supply is made of copper. If you follow the circuit through the interface you will also see that the negative (-) lead from the power supply is indirectly attached to the stainless steel electrode under the gas measuring tube. Since electrons are supplied at this electrode, this is where water will be reduced to hydrogen gas. Also, none of the metal “hook” of the electrode should be exposed outside the mouth of the graduated cylinder or else copper ions may plate onto it. In order to minimize the amount of water used for the electrolysis, begin by filling the 25 mL graduated cylinder and 100 mL beaker in a bucket of water (your instructor will demonstrate). Once the apparatus is clamped in place [you can use the clamp to help hold the stainless steel electrode in the proper position, centered in the cylinder mouth with no metal exposed below the cylinder], use a turkey baster to remove as much water as possible (you should have no more than about 60 mL in the beaker). Then add the 3 mL of acid using the hooked beral pipet (your instructor will demonstrate). 2. Measuring current again You can use the same technique to measure current in this experiment as you used in plating. Hook up the entire circuit but leave the copper electrode out of the solution. When you are ready to begin, wait until time zero appears on your calculator screen and then insert the copper electrode into the solution. Recording every 15 seconds will be sufficient. Watch the gas collection carefully. You don't want to exceed the 25 mL mark so be sure to pull the copper electrode out at the nearest time reading that will give close to 25 mL of gas. 3. Determining the gas volume in the graduated cylinder If you want to figure out how much hydrogen you have in the gas measuring tube, the volume will not be enough information. You will also need the temperature and pressure. You may remember a trick we used earlier in the year to make the pressure in a gas measuring tube equal to the room pressure. Be sure the liquid level inside the cylinder is the same as the liquid level in the beaker before recording the volume of gas. The room temperature and pressure will be posted. 4. The solution Once the electrolysis is complete, and the volume of gas has been measured, you need to drain the remaining liquid in the cylinder into the beaker (hence the limitation on total volume of water). The beaker of solution should then be transferred into a 100 mL cylinder and small rinsings of the beaker used to bring the total volume to 100.0 mL. Then pour the resulting solution into a clean 250 mL beaker and mix well. All of your copper(II) is now in this beaker. 1.5 mL of this solution is removed by pipet and added to 1.5 mL of 6 M NH3 in a cuvette to make the solution for the colorimeter (the blank is water). A equal volume of the standard copper(II) solution is treated in the same way for a reference solution. The remaining 98.5 mL is used for the titration. 5. The titration Once the solid KI is added, the solution turns brown and Na2S2O3 is added from the buret until the color becomes beige (BUT NOT WHITE). Then add 3 mL starch and titrate to white. Analysis 1. Determine the approximate number of moles of copper oxidized from the mass change of the electrode. (For this

experiment, you only have enough information to report an answer to ONE sig. fig.!!!) 2. Use your data to determine the moles of hydrogen gas collected during the electrolysis. Because the graduated

cylinder was intended to be read upright, there is a consistent error in reading the volume upside down. Repeated measurements show that the error is about +4%. For a 25 mL cylinder, this means that the volume you record looks 1 mL larger than it really is. Also, don't forget to correct for the water vapor pressure (table below).

 

 

Unit 12 36  

Vapor pressure of water

Temperature, oC Pressure, mm Hg Pressure, kPa

19

16.5

2.20

20

17.5 2.33

21

18.7 2.49

22

19.8 2.64

23

21.1 2.81

24

22.4 2.99

25

23.8 3.17

26 25.2 3.36

3. Compare the approximate moles of copper with the moles of hydrogen. Based on the balanced electrolysis

equation, do these values roughly agree? 4. Use your time and current data to determine the moles of Cu that dissolved. Compare this to #2. Is there good

agreement? 5. Based on the ratio of S2O3

2- to Cu2+ in the titration balanced equation and the volume of thiosulfate used in the titration, determine the moles of copper present in the 100 mL of solution. Remember, you only titrated 98.5 mL of the original solution [1.5 mL were removed]. Be sure to take this into account in your calculation. Compare this to #2. Is there good agreement?

6. Finally, determine the concentration of copper(II) ion in your mixture. Recall Beer’s Law which states that

absorbance is directly proportional to concentration and use the absorbance of the standard copper solution (treated with ammonia) as a comparison for your sample. Finally, use the concentration you determine to find the moles of copper(II) that must have been in your 100 mL sample. Is there good agreement with the result in #2?

Unit 12 37  

Unit 12 Sample Test The test will follow the usual format, with 5 multiple choice questions, three required problems, a choice of two out of four balanced net-ionic equations to write, and one essay question. The following are representative of typical multiple choice questions but do not necessarily indicate topics to be addressed on your actual test. _____ 1. 2-propanol, C3H7OH, can be oxidized to acetone, C3H6O, under acidic conditions. The two balanced half-reaction for this process are given below: C3H7OH C3H6O + 2 e- + 2 H+

6 H+ + 3 e- + CrO3 Cr3+ + 3 H2O In this reaction, 4 moles of CrO3 would react exactly to completion with a. 4 moles of 2-propanol d. 3 moles of 2-propanol b. 2 moles of 2-propanol e. 1 mole of 2-propanol c. 6 moles of 2-propanol _____ 2. Consider the reaction: C + H2O CO + H2

Which of the following statements is FALSE? a. carbon is oxidized in the reaction b. oxygen is reduced in the reaction c. carbon monoxide is the name of one product of this reaction d. the oxidation number of the product hydrogen is 0 _____3. During the electrolysis of a dilute aqueous solution of CoBr2, which substance is produced at the anode? a. Co b. O2 c. Br2 d. H2

_____ 4. Which of the following is likely to oxidize the chloride ion to the chlorine molecule? a. Zn2+ b. Fe3+ c. Na d. MnO4

- e. Mn2+

_____5. In a galvanic cell consisting of zinc and silver in appropriate electrolytes, which of the following is TRUE? a. the zinc electrode gains mass and the silver electrode loses mass b. electrons flow from the silver electrode to the zinc electrode c. cations in the electrolyte move toward the silver electrode d. the voltage remains constant as the cell operates

Unit 12 38  

The next section consists of representative problems which might be found in the problems section. 6. A steady current of 1.00 amp is passed through an electrolytic cell containing a 0.10 M solution of CuF2 using inert graphite electrodes until 1.54 g of copper is deposited. ___________a. At which electrode is the copper deposited? (anode/cathode) ___________b. How many minutes does the current flow to obtain this deposit? ___________ c. How many Litres of what gas would be produced at the other electrode while the copper is deposited? (assume STP) 7. Consider the reaction: Cu2+(aq) + Pb(s) Cu(s) + Pb2+(aq)

a. Draw and label a galvanic cell for this reaction, including the anode, cathode, flow of electrons and direction of flow for the anions and cations. b. Calculate E o

c. Calculate ΔGo

Unit 12 39  

8. Balance the following redox reaction and identify the oxidizing agent, reducing agent, the substance oxidized and the substance reduced. SO3

2- + CrO42- SO4

2- + Cr(OH)3 (OH-)

9. For each of the following pairs, circle the metal which would be protected from corrosion when placed in contact with the other metal: a. Zn and Cu b. Mg and Fe c. Sn and Pb d. Sn and Cu

Unit 12 40  

The next section consists of representative reactions to complete and write balanced net-ionic equations for. Note that some reactions do not occur in aqueous solution and thus molecular equations are all that would be needed. Each student is expected to choose two from this section. Phase symbols ((s), (aq) etc.) are required for full credit.

10. dihydrogen sulfide gas is bubbled into an acidic solution of sulfate ions; solid sulfur and sulfur dioxide gas are among the products ______________________________________________________________________________ ______________________________________________________________________________ ______________________________________________________________________________ 11. copper(II) chloride + lead(II) nitrate ___________________________________________ ______________________________________________________________________________ ______________________________________________________________________________ type: _________________ 12. zinc metal + sulfuric acid ____________________________________________________ ______________________________________________________________________________ ______________________________________________________________________________ type: _________________ 13. nitric acid + calcium nitrate __________________________________________________ ______________________________________________________________________________ ______________________________________________________________________________ type: _________________ The final section of the test will consist of one essay question selected from the following topics: --- use of Standard Reduction Potential Table to predict the products of electrolysis --- the use and function of starch indicator --- sacrificial metals in corrosion prevention --- rechargeable batteries