COPPER-SILVER REACTION Chemistry 121, J. Chen Chemical reactions can be divided into several categories. Some chemical reactions fall into more than one category. For example, let’s take a look at the chemical reaction that we will be investigating in this experiment:

2AgNO3(aq) + Cu(s) -> 2Ag(s) + Cu(NO3)2(aq) This reaction is an example of a SINGLE REPLACEMENT reaction, in which a lone element (solid copper in this case) will "replace" an element in a compound (silver nitrate in this case) to produce a different lone element and a different compound. This reaction also is an example of a REDOX reaction (short for ”reduction-oxidation”). In a redox reaction, one of the elements on the reactant side will LOSE at least one electron to produce the products, while another element on the reactant side will GAIN at least one electron to produce the products. The reactant element that has LOST electrons has been OXIDIZED, while the reactant element that has GAINED electrons has been REDUCED. The above balanced chemical equation declares that every TWO moles of aqueous silver nitrate must react with ONE mole of solid copper metal in order to produce two moles of sold silver metal and one mole of aqueous copper(II) nitrate. Thus, in order to know how much of the products we can possibly make (the THEORETICAL YIELD), we must know how many moles of each reactant are being used in the reaction. We can calculate the number of moles of reactants being used by WEIGHING our reactants and converting the masses to moles by using the MOLAR MASSES of the reactants as conversion factors. We can also determine how close we are to obtaining the theoretical yield by WEIGHING the final product (the EXPERIMENTAL YIELD) and calculating what percentage of the theoretical yield that this represents (the PERCENT YIELD). IMPORTANT NOTES BEFORE STARTING THE EXPERIMENT: 1. We will be using silver nitrate in this experiment. Please avoid getting the siiver nitrate on your hands since silver nitrate will stain your skin. The quantity we will be using will not harm your skin, and the staining will subside within a few days. 2. In order to obtain an accurate value for the experimental yieid, you will place your product in the oven before you leave, and allow the product to dry overnight. You must come back the following day to obtain the final mass value needed for your calculations.

PROCEDURE: Obtain a vial containing silver nitrate. Weigh the vial containing the silver nitrate and record the mass. Also record the balance number being used, and USE THE SAME BALANCE FOR THE REST OF THE EXPERIMENT. In a 25 mL graduated cylinder, measure about 20 mL of distilled water. Empty the silver nitrate from the vial into the graduated cylinder, then stir carefully with a glass rod until the silver nitrate is dissolved. Weigh the empty vial, and record the mass. The mass of the “full vial” minus the mass of the “empty vial” will equal the mass of silver nitrate used in this reaction. Obtain a piece of copper wire about one foot long. Coil about ¾ of the wire around a glass rod or pen,, then stretch the wire so that it is about 1 cm taller than the graduated cylinder. Weigh the wire and record the mass. Place the wire (coiled end down) into the graduated cylinder, and leave it alone for the next 15 minutes. During this time, observe what is happening in the graduated cylinder, and record your observations. Also during this time, clean and dry a 100 mL beaker, to be used as the container for your final product. Use a market to mark the glass with your initials or some other mark that will allow you to identify which beaker is yours. Then, weigh the empty beaker and record the mass. After the 15 minutes are up, gently shake the product free from the wire, and allow the reaction to continue for 5 minutes more, to allow the reaction to go to completion. Once the reaction is complete, remove the wire from the graduated cylinder, and squirt the wire with distilled water to rinse any product crystals still stuck to the wire back into the graduated cylinder. Then, dip the wire into a test tube of acetone (in the hood), and allow the wire to dry naturally. Once the wire is dry, weigh the wire and record the mass. Carefully decant (pour off) most of the blue solution from the graduated cylinder into a large waste beaker. Then, use your wash bottle containing distilled water to transfer the crystals into your clean, dry, and weighed 100 mL beaker. Throughout this process, do your best not to lose any of the crystals. After the crystals are transferred to the beaker, wash the crystals by adding more distilled water to the beaker, letting the crystals settle, and decanting the liquid. Repeat the process until the blue color of the solution has completely disappeared. A total of 3 rinses should do the trick. Again, do your best not to lose any of the crystals!

Once you have gotten as much of the liquid out of the beaker as you can without losing crystals, place the beaker in to the oven and allow the sample to dry overnight. Clean, dry, and return the equipment. Pour all of your chemical waste into the labeled waste container in the hood. The following day, retrieve your sample (it may already have been taken out of the oven for you). If the beaker is warm to the touch, then let it cool down for about 5 minutes. Then, use the same balance that you used the previous day to weigh your beaker with the sample, and record the mass. Transfer your product to the labeled waste container in the hood. DATA AND OBSERVATIONS Balance Number ____________________________ Partner’s Name ____________________________ ID on Product Beaker _____________________ Mass Before

(grams) Mass After

(grams)

Vial of AgNO3

Copper Wire

Product Beaker

Observations to include: Colors, times, what the products look like, and anything else that you feel may be worth noting!

CALCULATIONS (Show all values and a set-­up for each calculation):

1. Mass of AgNO3 used in reaction

2. Mass of copper wire used in reaction

3. Mass of silver metal produced in reaction

4. Moles of AgNO3 used in reaction

5. Moles of copper wire used in reaction

6. Moles of silver metal produced in reaction

7. Ratio of moles Ag produced to moles AgNO3 used (round this ratio off to

the nearest whole number)

8. Ratio of moles Ag produced to moles Cu used (round this ratio off to the

nearest whole number)

9. Number of atoms of Ag(s) produced

10. Theoretical yield of Ag(s) in grams, based on the amount of AgNO3 used

11. Percent yield of Ag(s)

DISCUSSION QUESTIONS:

1. Looking at the balanced chemical equation, please identify:

(a) The atom or ion on the reactant side that has been OXIDIZED to form the

products. (explain why?)

(b) The atom or ion on the reactant side that has been REDUCED to form the

products. (explain why?)

2. Even though the reaction was carried to completion, why wasn't all the

copper wire used up?

3. For each of the following, EXPLAIN in two or three sentences whether the

experimental mole ratio of Ag to Cu would be HIGHER THAN, LOWER

THAN, or EQUAL TO the theoretical mole ratio.

(a) Some AgNO3 crystals were left in the vial after the AgNO3 was emptied

into the graduated cylinder.

(b) Some of the solid silver crystals remained stuck to the copper wire after

the reaction and were not transferred to the product beaker.

(c) The solid silver crystals were not rinsed completely clean before being

dried and weighed.

4. If the mass of copper wire used in your reaction was only 0.012 grams, but

the amount of silver nitrate used was unchanged, then how many grams of

silver metal could be produced? What is the limiting reactant in this case?

Conclusion:

As always, conclude with one or two paragraphs that discuss:

a) What was learned in the experiment;

b) Possible sources of error; and

c) What could be done to improve the experiment.

Even if the experiment went “perfectly,” please come up with some possible

answers for (b) and (c) as if the experiment did not go “perfectly.” “Human

error” and “errors in calculations” are NOT acceptable reasons for an

imperfect experiment! Please come up with OTHER possible reasons.

Experiment 7 65

Experiment 7Determination of the formula of a metal oxide

Background

Through the use of chemical symbols and numerical subscripts, the formula of a compoundcan be written. The simplest formula that may be written is the empirical formula. In thisformula, the subscripts are in the form of the simplest whole number ratio of the atoms ina molecule or of the ions in a formula unit. The molecular formula, however, representsthe actual number of atoms in a molecule. For example, although CH2O represents theempirical formula of the sugar, glucose, C6H12O6, represents the molecular formula. Forwater, H2O, and carbon dioxide, CO2, the empirical and the molecular formulas are thesame. Ionic compounds are generally written as empirical formulas only; for example,common table salt is NaCl.

The formation of a compound from pure components is independent of the source ofthe material or of the method of preparation. If elements chemically react to form acompound, they always combine in definite proportions by weight. This concept is knownas the Law of Constant Composition.

If the weight of each element that combines in an experiment is known, then thenumber of moles of each element can be determined. The empirical formula of thecompound formed is the ratio between the number of moles of elements in the compound.This can be illustrated by the following example. If 32.06 grams of sulfur is burned in thepresence of 32.00 grams of oxygen, then 64.06 grams of sulfur dioxide results. Thus

and the mole ratio of sulfur � oxygen is 1 : 2. The empirical formula of sulfur dioxide isSO2. This also is the molecular formula.

In this experiment, the moderately reactive metal, magnesium, is combined withoxygen. The oxide, magnesium oxide, is formed. The equation for this reaction, based onthe known chemical behavior, is

If the mass of the magnesium is known and the mass of the oxide is found in theexperiment, the mass of the oxygen in the oxide can be calculated:

As soon as the masses are known, the moles of each component can be calculated. The molescan then be expressed in a simple whole number ratio and an empirical formula written.

mass of magnesium oxide� mass of magnesium mass of oxygen

2Mg(s) � O2(g) heat¶¶l 2MgO(s)

32.00 g O

16.00 g/mole O � 2 moles of oxygen

32.06 g S

32.06 g/mole S � 1 mole of sulfur

In the present experiment, magnesium metal is heated in air. Air is composed ofapproximately 78% nitrogen and 21% oxygen. A side reaction occurs between some of themagnesium and the nitrogen gas:

Not all of the magnesium is converted into magnesium oxide; some becomesmagnesium nitride. However, the magnesium nitride can be converted to magnesiumoxide by the addition of water:

As a result, all of the magnesium is transformed into magnesium oxide.

Mg3N2(s) � 3H2O(l) heat¶¶l 3MgO(s) � 2NH3(g)

3Mg(s) � N2(g) heat¶¶l Mg3N2(s)

66 Experiment 7 .

� 4.03 g� 2.43 g� 1.60 g

EXAMPLE

When 2.43 g of magnesium was burned in oxygen, 4.03 g of magnesium oxidewas produced.

No. of moles of magnesium � � 0.100 moles

No. of moles of oxygen � � 0.100 moles

The molar ratio is 0.100 � 0.100 � 1 � 1The empirical formula is Mg1 O1 or MgO.

%Mg � � 100 � 60.3%2.43 g4.03 g

1.60 g16.00 g/mole

2.43 g24.31 g/mole

mass of magnesium oxide� mass of magnesium mass of oxygen

Objectives

1. To prepare a metal oxide.2. To verify the empirical formula of a metal oxide.3. To demonstrate the Law of Constant Composition.

Procedure

CAUTION!

A hot crucible can cause severe burns if handled improperly. Be sure to allow thecrucible to cool sufficiently before handling. Always handle a hot crucible withcrucible tongs.

Cleaning the Crucible

1. Obtain a porcelain crucible and cover. Carefully clean the crucible in the hood byadding 10 mL of 6 M HCl to the crucible; allow the crucible to stand for 5 min. with theacid. With crucible tongs, pick up the crucible, discard the HCl, and rinse the cruciblewith distilled water from a plastic squeeze bottle.

2. Place the crucible in a clay triangle, which is mounted on an iron ring and attached to aring stand. Be sure the crucible is firmly in place in the triangle. Place the cruciblecover on the crucible slightly ajar (Fig. 7.1a).

Experiment 7 67

Figure 7.1(a) Heating the crucible. (b) Picking up the cruciblewith crucible tongs.

b

a

3. Begin to heat the crucible with the aid of a Bunsen burner in order to evaporate water.Increase the heat, and, with the most intense flame (the tip of the inner blue cone), heatthe crucible and cover for 5 min.; a cherry red color should appear when the bottom isheated strongly. Remove the flame. With tongs, remove the crucible to a heat-resistantsurface and allow the crucible and cover to reach room temperature.

4. When cool, weigh the crucible and cover to 0.001 g (1). (Be sure to handle with tongssince fingerprints leave a residue.)

5. Place the crucible and cover in the clay triangle again. Reheat the crucible to the cherryred color for 5 min. Allow the crucible and cover to cool to room temperature. Reweighwhen cool (2). Compare weight (1) and weight (2). If the weight differs by more than0.005 g, heat the crucible and cover again for 5 min. and reweigh when cool. Continueheating, cooling, and weighing until the weight of the crucible and cover are constant towithin 0.005 g.

Forming the Oxide

1. Using forceps to handle the magnesium ribbon, cut a piece approximately 12 cm inlength and fold the metal into a ball; transfer to the crucible. Weigh the crucible, cover,and magnesium to 0.001 g (3). Determine the weight of magnesium metal (4) bysubtraction.

2. Transfer the crucible to the clay triangle; the cover should be slightly ajar (Fig. 7.1a).

3. Using a small flame, gently apply heat to the crucible. Should fumes begin to appear,remove the heat and cover the crucible immediately. Again place the cover ajar andcontinue to gently heat for 10 min. (If fumes appear, cover as before.) Remove the flameand allow the assembly to cool for 2 min. With tongs, remove the cover. If themagnesium has been fully oxidized, the contents should be a dull gray. Shiny metalmeans there is still free metal present. The cover should be replaced as before and thecrucible heated for an additional 5 min. Reexamine the metal and continue heatinguntil no shiny metal surfaces are present.

4. When all the metal appears as the dull gray oxide, half-cover the crucible and gently heatwith a small Bunsen flame. Over a period of 5 min., gradually adjust the intensity of theflame until it is at its hottest, then heat the crucible to the cherry red color for 5 min.

Completing the Reaction

1. Discontinue heating and allow the crucible assembly to cool to room temperature.Remove the cover and, with a glass rod, carefully break up the solid in the crucible.With 0.5 mL (10 drops) of distilled water dispensed from an eye dropper, wash the glassrod, adding the water to the crucible.

2. Set the cover ajar on the crucible and gently heat with a small Bunsen flame toevaporate the water. (Be careful to avoid spattering while heating; if spattering occurs,remove the heat and quickly cover the crucible completely.)

3. When all the water has been evaporated, half-cover the crucible and gradually increaseto the hottest flame. Heat the crucible and the contents with the hottest flame for 10 min.

4. Allow the crucible assembly to cool to room temperature. Weigh the cool crucible, cover,and magnesium oxide to 0.001 g (5).

5. Return the crucible, cover, and magnesium oxide to the clay triangle. Heat at full heatof the Bunsen flame for 5 min. Allow to cool and then reweigh (6). The two weights, (5)and (6), must agree to within 0.005 g; if not, the crucible assembly must be heated for 5min., cooled, and reweighed until two successive weights are within 0.005 g.

Calculations

1. Determine the weight of magnesium oxide (7) by subtraction.

2. Determine the weight of oxygen (8) by subtraction.

3. From the data obtained in the experiment, calculate the empirical formula ofmagnesium oxide.

68 Experiment 7 .

Chemicals and Equipment

1. Clay triangle2. Porcelain crucible and cover3. Crucible tongs4. Magnesium ribbon5. Eye dropper6. 6 M HCl

Experiment 7 71

Experiment 7

REPORT SHEET

1. Weight of crucible and cover (1) ______________ g

2. Weight of crucible and cover (2) ______________ g

3. Weight of crucible, cover, and Mg (3) ______________ g

4. Weight of Mg metal (4): (3) � (2) ______________ g

5. Weight of crucible, cover, and oxide (5) ______________ g

6. Weight of crucible, cover, and oxide (6) ______________ g

7. Weight of magnesium oxide (7): (6) � (2) ______________ g

8. Weight of oxygen (8): (7) � (4) ______________ g

9. Number of moles of magnesium (4)/24.31 g/mole ______________ moles

10. Number of moles of oxygen (8)/16.00 g/mole ______________ moles

11. Simplest whole number ratio of Mg atoms to O atoms � (9)�(10) _________ : _________

12. Empirical formula for magnesium oxide ______________

13. % Mg in the oxide from data % � [(4)/(7)] � 100 ______________ %

14. % Mg calculated from the formula MgO % � [24.31 g/40.31 g] � 100 ______________ %

15. Error

% � � 100 ______________ %(14) � (13)(14)

NAME SECTION DATE

PARTNER GRADE

POST-LAB QUESTIONS

1. Write the balanced equation for the formation of iron(II) oxide, FeO, from the elementsof iron, Fe, and oxygen, O2.

2. Write the two chemical equations that describe the conversion of magnesium intomagnesium oxide.

3. What error in calculation would result if, in the procedure for forming the magnesiumoxide, the fumes in the initial heating were allowed to escape?

4. Mercuric oxide decomposes when heated and forms free mercury metal and releasesoxygen gas. If a sample of the oxide, 2.096 g, is decomposed, 1.942 g of mercury isobtained. Determine the empirical formula for mercuric oxide. (Hint: first determinethe number of moles of mercury and oxygen in the sample.)

72 Experiment 7 .