on the same layer
on layer above
TOTAL CN of the interior sphere
c. Sphere packing that has this number (write below) of adjacent and touching nearest neighbors is referred to as close-packed. Non-close-packed structures will have lower coordination numbers.
d. Are the two unit cells the identical?
e. If they are the same, how are they related? If they are different, what makes them different?
f. Is the face-centered cubic unit cell aba or abc layering? Draw a z-diagram.
III.Interstitial sites and coordination number (CN)
a. If the spheres are assumed to be ions, which of the spheres is most likely to be the anion and which the cation, the colorless spheres or the colored spheres? Why?
b. Consider interstitial sites created by spheres of the same size. Rank the interstitial sites, as identified by their coordination numbers, in order of increasing size (for example, which is
biggest, the site with coordination number 4, 6 or 8?).
c. Using basic principles of geometry and assuming that the colorless spheres are the same anion with radius r A in all three cases, calculate in terms of rA the maximum radius, rC, of the cation that will fit inside a hole of CN 4, CN 6 and CN 8. Do this by calculating the ratio of the radius of to cation to the radius of the anion: rC/ rA.
d. What terms are used to describe the shapes (coordination) of the interstitial sites of CN 4, CN 6
and CN 8?
CN 4: ________________
CN 6: _______________
CN 8: ________________
IV.Ionic Solids
A. Cesium Chloride
1. Fill the table below for Cesium Chloride
Table 7.8.
Colorless spheres Green spheres
Type of cubic structure
Atom represented
2. Using the simplest unit cell described by the colorless spheres, how many net colorless and net green spheres are contained within that unit cell?
3. Do the same for a unit cell bounded by green spheres as you did for the colorless spheres in
question 4.
4. What is the ion-to-ion ratio of cesium to chloride in the simplest unit cell which contains both?
(Does it make sense? Do the charges agree?)
B. Calcium Fluoride
1. Draw the z diagrams for the layers (include both colorless and green spheres).
2. Fill the table below for Calcium Fluoride
Table 7.9.
Colorless spheres Green spheres
Type of cubic structure
Atom represented
3. What is the formula for fluorite (calcium fluoride)?
C. Lithium Nitride
1. Draw the z diagrams for the atom layers which you have constructed.
2. What is the stoichiometric ratio of green to blue spheres?
3. Now consider that one sphere represents lithium and the other nitrogen. What is the formula?
4. How does this formula correspond to what might be predicted by the Periodic Table?
D. Zinc Blende and Wurtzite
Fill in the table below:
Table 7.10.
Zinc Blende Wurtzite
Stoichiometric ratio of colorless to pink spheres
Formula unit (one sphere represents and the other the sulfide ion)
Compare to predicted from periodic table
Type of unit cell
Solutions
Chapter 8. Colorful Copper
Experiment 3: Colorful Copper
Objective
To observe, describe and explain the products of a number of chemical reactions of the
transition metal copper.
To use several techniques in recovering copper from solution.
To understand the concept of percent yield.
Grading
Pre-lab (10%)
Lab Report (80%)
TA points (10%)
Before Coming to Lab…
Read the lab instructions
Complete the pre-lab, due at the beginning of the lab.
Introduction
Copper is a soft metal with a characteristic color that we often call "copper-colored", a bright orange-brown color. Copper is relatively inert chemically; it does not readily oxidize (react with oxygen) in air and is react when exposed to simple mineral acids such as sulfuric or hydrochloric acid. One of the most popular uses of copper is in the computer industry where it is used to build the integrated circuits and chips. It is beginning to replace aluminum for this application due to the resulting decrease in costs. Copper is also good at conducting electricity because it has so many free electrons that allow for the efficient flow of current.
In this lab you will preform a series of reactions with copper and observe a variety of distinctive and colorful products. Most chemical syntheses involve the separation and then purification of a desired product from unwanted side products. The common methods of separation are filtration,
sedimentation, decantation, extraction, chromatography and sublimation.
This experiment is designed as a quantitative evaluation of your laboratory skills in carrying out a series of chemical reactions, purification and analyses with copper. You will employ two
fundamental types of chemical reactions, namely oxidation-reduction (redox) and metathesis
(exchange) reactions to recover pure copper with maximum efficiency. The chemical reactions
involved are the following.
Redox: Cu( s)+4HNO3(aq)→Cu(NO3)2(aq)+NO2( g)+2H2 O( l)[1]
*Metathesis: Cu(NO3)2(aq)+2NaOH(aq)→Cu(OH)2( s)+2NaHO3(aq)[2]
Dehydration: Cu(OH)2( s)+heat→CuO( s)+ H 2 O( g)[3]
Metathesis: CuO( s)+ H 2SO4(aq)→CuSO4(aq)+ H 2 O( l)[4]
Redox: 3CuSO4(aq)+2Al( s)→Al2(SO4)3(aq)+Cu( s)[5]
Each of these reactions proceeds to completion and in the case of a metathesis reaction,
completion is reached when one of the components is removed from the solution in form of a gas
or an insoluble precipitate. This is the case for reactions [1], [2], and [3]. In reactions [1] and [3] a gas is formed and in reaction [2] an insoluble precipitate is formed (Reaction [5] proceed to
completion because copper is more difficult to oxidize than aluminum).
Metathesis (Exchange) Reaction Defined in Chapter 4 of your textbook: 'One of the following is
needed to drive a metathesis reaction: the formation of a precipitate, the generation of a gas, the production of a weak electrolyte, or the production of a nonelectrolyte.'
Oxidation-Reduction (Redox) Reactions. This involves the loss of electrons from one components
and an addition of electrons to the other component as the reaction proceeds (the are transferred from one atom to another). The component that loses electrons is said to be oxidized; the one that gains electrons is then reduced. Such reactions are important for the production of electricity due to the energy produced from an electron transfer.
The percent yield of the copper can be expressed as the ratio of the recovered mass to initial mass, multiplied by 100:
% yield = (recovered mass of Cu/initial mass of Cu) x 100
Experimental Procedure
SAFETY PRECAUTIONS: Wear safety glasses and gloves when handling acids. Work in the fume
hood. Acetone and Methanol are Flammable, so keep them away from flames. Avoid breathing
any vapors, especially methanol, as it is very toxic.
1. Preheat a hot plate in the fume hood.
2. Place one strip of fine copper wire (Remember to record the actual weight, approximately 0.5 g) in a 250 mL beaker. Bend the copper wire so that it rests on the bottom of the beaker.
3. Slowly add 30 mL of 6 M HNO3. Perform this step in the fume hood, as the gas produced is
toxic. You may need to gently heat the solution. Observe the color of the gas and solution. Place a wet paper towel and watchglass over the beaker to dissolve the gas. Ask your TA if you are unsure how to do this.
4. Wait until the copper wire dissolves completely before proceeding.
5. Remove beaker from hot plate with a hot hand. Remove watchglass and paper towel. Be sure
that the hood is closed as much as possible at this point to avoid breathing in the gas.
6. Add 100 mL of deionized water (Do not use regular water!)
7. Slowly add 30 mL of 6 M NaOH. Note the color of the precipitate and evolution of heat.
8. Add 2 or 3 boiling chips and carefully heat the solution just to boiling point. Note any color change. While waiting for this solution to heat, begin heating ~200 mL deionized water (in a 400
mL beaker) for the next step.
9. Decant some of the supernatant liquid into a 500 mL beaker (Try not to lose any solid while
decanting. It is ok to leave some of the liquid behind.). Add about 200 mL of very hot deionized water and allow the precipitate to settle. Decant once more. What are you removing by washing
and decantation?
10. Heat the beaker containing the solid for 20 minutes to reduce the volume of the solution (or until the volume has been reduced by half). Having a more concentrated solution will make the
following steps proceed faster.
11. Add 15 mL of 6 M H 2SO4 to the black tarry substance in the beaker (not the solution that you have decanted) while stirring with a glass rod. What copper compound is present now?
12. Remove the boiling chips.
13. In the hood, add 5-10 one-inch squares of aluminum foil and 5-10 ml of concentrated HCl,
noting any color changes.STIR WELL. An ideal ratio should be 7 pieces of Aluminum to 10 ml of
acid but you might need more or less depending on the success of your previous steps. If your
solution turns green stop and ask your TA for help. Otherwise, continue to add pieces of aluminum until the supernatant is not blue. Identify what forms on the surface of the aluminum. What is
present in the solution? What gas is formed in the reaction? How do you know?
14. When gas evolution has ceased, decant the solution and transfer the precipitate to a
preweighed 100 mL beaker and record its mass on your report form. Wash the precipitated copper
with about 5 mL of distilled water and allow it to settle before you decant the solution, and repeat the process. Then wash the precipitate with about 5 mL of methanol. Allow the precipitate to
settle, and then decant the methanol. Finally, wash the precipitate with about 5 mL of acetone.
Allow the precipitate to settle again and then decant the acetone from the precipitate. What are you removing by washing?
15. Then use the microwaves to dry your product. Start with 30 seconds and then try 10 seconds
more until it is dry. Be sure to let your product cool before weighing it. Then calculate the final mass of copper and from there the weight and percent yield can be determined. Compare the mass
with your initial mass and calculate the percent yield. What color is your copper sample in the
final step?
Is it uniform in appearance?
Suggest possible sources of error in this experiment.
Solutions
Chapter 9. Metathesis: To Exchange or Not?
Lab 4: Metathesis: To Exchange or Not
Objectives
To give practice writing equations for metathesis reactions, including net ionic equations
To illustrate the concept of solubility and the effect of temperature and crystallization
Grading
You will be determined according to the following:
Pre-lab (10%)
Must attach graph
Lab Report Form (80%)
Must include detailed observations for each reaction
TA Evaluation of lab procedure (10%)
Before Coming to Lab…
Complete the pre-lab exercise, including the plot (due at the beginning of lab)
Read the instructions for the lab and refresh your memory on anything that isn’t clear by
reading your textbook
Introduction
In molecular equations for many aqueous reactions, cations and anions appear to exchange
partners. These reactions conform to the following general equation:
Equation 1: AX+BY→AY+BX
These reactions are known as metathesis reactions. For a metathesis reaction to lead to a net
change in solution, ions must be removed from the solution. In general, three chemical processes can lead to the removal of ions from solution, comcomitantly serving as a driving force for
metathesis to occur:
1. The formation of a precipitate2. The formation of a weak electrolyte or nonelectrolyte3. The
formation of a gas that escapes from solution
The reaction of barium chloride with silver nitrate is a typical example:
Equation 2: BaCl2(aq)+2AgNO3(aq)→Ba(NO3)2(aq)+2AgCl( s)
This form of the equation for this reaction is referred to as the molecular equations. Since we
know that the salts BaCl2, AgNO3, and Ba(NO3)2 are strong electrolytes and are completely
dissociated in solution, we can more realistically write the equation as follows:
Equation 3: Ba2+(aq)+2Cl−(aq)+2Ag+(aq)+2NO3−(aq)→Ba2+(aq)+2NO3−(aq)+2AgCl( s)
This form, in which all ions are shown, is known as the complete ionic equation. Reaction occurs because the insoluble substance AgCl precipitates out of solution. The other product, barium
nitrate, is soluble in water and remains in solution. We see that Ba2+ and NO3− ions appear on both sides of the equation and thus do not enter into the reaction. Such ions are called spectator ions. If we eliminate or omit them from both sides, we obtain the net ionic equation:
Equation 4: Ag+(aq)+Cl−(aq)→AgCl( s)
This equation focuses our attention on the salient feature of the reaction: the formation of the precipitate AgCl. It tells us that solutions of any soluble Ag+saltand any soluble Cl−salt, when mixed, will form insoluble AgCl. When writing net ionic equations, remember that only strong
electrolytes are written in the ionic form. Solids, gases, nonelectrolytes, and weak electrolytes are written in the molecular form. Frequently the symbol (aq) is omitted from ionic equations. The
symbols (g) for gas and (s) for solid should not be omitted. Thus, Equation 4 can be written as
Equation 5: Ag++Cl−→AgCl( s)
Consider mixing solutions of KCl and NaNO3. The ionic equation for the reaction is
Equation 6: K+(aq)+Cl−(aq)+Na+(aq)+NO3−(aq)→ K+(aq)+NO3−(aq)+Na+(aq)+Cl−(aq)
Because all the compounds are water-soluble and are strong electrolytes, they have been written in the ionic form. They completely dissolve in water. If we eliminate spectator ions from the
equation, nothing remains. Hence, there is no reaction: Equation 7:
K+(aq)+Cl−(aq)+Na+(aq)+NO3−(aq)→no reaction
Metathesis reactions occur when a precipitate, a gas, a weak electrolyte, or a nonelectrolyte is formed. The following equations are further illustrations of such processes.
Formation of a Gas
Molecular equation: Equation 8: 2HCl(aq)+Na2 S(aq)→2NaCl(aq)+ H 2 S( g) Complete ionic equation: 2H+(aq)+2Cl−(aq)+2Na+(aq)+ S 2−(aq)→2Na+(aq)+2Cl−(aq)+ H 2 S( g) Net ionic equation: 2H+(aq)+ S 2−(aq)→ H 2 S( g)
or
2H + + S 2 − → H 2 S ( g )
Formation of a Weak Electrolyte
Molecular equation:
HNO 3 ( aq ) + NaOH ( aq ) → H 2 O ( l ) + NaNO 3 ( aq )
Complete ionic equation:
H + ( aq ) + NO 3 − ( aq ) + Na + ( aq ) + OH − ( aq ) → H 2 O ( l ) + Na + ( aq ) NO 3 − ( aq ) Net ionic equation:
H + ( aq ) + OH − ( aq ) → H 2 O ( l )
In order to decide if a reaction occurs, we need to be able to determine whether or not a
precipitate, a gas, a nonelectrolyte, or a weak electrolyte will be formed. The following brief
discussion is intended to aid you in this regard. Table 1 summarizes solubility rules and should be consulted while performing this experiment.
The common gases are CO2, SO2, H 2 S, and NH3. Carbon dioxide and sulfur dioxide may be regarded as resulting form the decomposition of their corresponding weak acids, which are
initially formed when carbonate and sulfite salts are treated with acid:
H 2 CO 3 ( aq ) → H 2 O ( l ) + CO 2 ( g )
and
H 2 SO 3 ( aq ) → H 2 O ( l ) + SO 2 ( g )
Ammonium salts form NH3 when they are treated with strong bases:
NH 4 + ( aq ) + OH − → NH 3 ( g ) + H 2 O ( l )
Table 1 Solubility Rules
Table 9.1.
Water-soluble salts
Na + ,K + ,NH 4 +
All sodium, potassium, and ammonium salts are soluble.
NO 3 − ,CIO 3 − ,C 2 H 3 O 2 − All nitrates, chlorates, and acetate are soluble.
Cl −
All chlorides are soluble except AgCl, Hg2Cl2, and
.
Br −
All bromides are soluble except AgBr, Hg2Br2,
, and
.
I −
All iodides are soluble except AgI, Hg2 I 2, PbI2, and HgI2.0
All sulfates are soluble except
, SrSO
SO
4, BaSO4, Hg2SO4,
4 2 −
PbSO4, and Ag2SO4.
Table 9.2.
Water-
insoluble salts
CO32−, SO32−,
All carbonates, sulfites, phosphates, and chromates are insoluble except those of
PO43−
CrO42− alkali metals and NH4+.
All hydroxides are insoluble except those of alkali metals and
,
OH −
, and Ba(OH)2.
All sulfides are insoluble except those of the alkali metals, alkaline earths, and
S 2 −
NH4+.
*Slightly soluble.
Table 2 Strong Electrolytes
Table 9.3.
Salts
All common soluble salts
Acids HClO4, HCl, HBr, HI, HNO3, and H 2SO4 are strong electrolytes; all others are weak.
Alkali metal hydroxides, Ca(OH)
Bases
2, Sr(OH)2, and Ba(OH)2 are strong electrolytes; all
others are weak.
Which are the weak electrolytes? The easiest way of answering this question is to identify all of the strong electrolytes, and if the substance does not fall in that category then it is a weak
electrolyte. Note, water is a nonelectrolyte. Strong electrolytes are summarized in Table.2.
In the first part of this experiment, you will study some metathesis reactions. In some instances it will be very evident that a reaction has occurred, whereas in others it will not be so apparent. In the doubtful case, use the guidelines above to decide whether or not a reaction has taken place.
You will be given the names of the compounds to use but not their formulas. This is being done
deliberately to give practice in writing formulas from names.
In the second part of this experiment, you will study the effect of temperature on solubility. The effect that temperature has on solubility varies from salt to salt. We conclude that mixing
solutions of KCl and NaNO3 resulted in no reaction (see Equations 6 and 7). What would happen if we cooled such a mixture? The solution would eventually become saturated with respect to one of
the salts, and crystals of that salt would begin to appear as its solubility was exceeded.
Examination of Equation 6 reveals that crystals of any of the following salts could appear
initially: KNO3, KCl, NaNO3, or NaCl.Consequently, if a solution containing Na+, K+, Cl−, and NO3− ions is evaporated at a given temperature, the solution becomes more and more concentrated
and will eventually become saturated with respect to one of the four compounds. If a evaporation is continued, that compound will crystallize out, removing its' ions from solution. The other ions will remain in solution and increase in concentration. Before beginning this laboratory exercise you are to plot a graph of the solubilities of the four salts given in Table 3 on your report sheet.
Experimental Procedure
Part 1: Metathesis Reactions
CAUTION WEAR EYE PROTECTION
1. The report sheet lists 16 pairs of chemicals that are to be mixed. Use about 1 mL of the
reagents to be combined as indicated on the report sheet.
2. Mix the solutions in small test tubes and record your observations on the report sheet. If there is no reaction, write N.R. (The reactions need not be carried out in the order listed. In order to reduce congestion at the reagent shelf, half the class will start in reverse order). Dispose of the contents of your test tubes in the designated receptacles.
Part 2: Solubility, Temperature and Crystallization
1. Place 8.5 g of sodium nitrate and 7.5 g of potassium chloride in a 100-mL beaker and add 25
mL of water. Warm the mixture on an hotplate, stirring, until the solids completely dissolve.
2. Assuming a volume of 25mL for the solution, calculate the molarity of the solution with
respect to NaNO3, KCl, NaCl, and KNO3, and record these molarities on your report form.
3. Cool the solution to about 10°C by placing the beaker in ice water in a 600-mL beaker and stir the solution carefully with a thermometer, being careful not to break it.
4. When no more crystals form, at approximately 10°C, filter the cold solution quickly and allow the filtrate to drain thoroughly into an evaporating dish. Dry the crystals between two dry
pieces of filter paper or paper towels.
5. Examine the crystals with a magnifying glass (or fill a Florence flask with water and look at the crystals through it). Describe the shape of the crystals—that is, needles, cubes, plates,
rhombs, and so forth on your report form.
6. Based upon your solubility graph, which compound crystallized out of solution and write that in the appropriate place on your report form
7. Evaporate the filtrate to about half of its volume using a Bunsen burner and ring stand. A
second crop of crystals should form. Record the temperature and rapidly filter the hot solution, collecting the filtrate in a clean 100-mL beaker.
8. Dry the second batch of crystals between two pieces of filter paper and examine their shape.
Compare their shape with the first batch of crystals.
9. Based upon your solubility graph, what is this substance?
10. Finally, cool the filtrate to 10°C while stirring carefully with a thermometer to obtain a third crop of crystals. Carefully observe their shapes and compare them with those of the first and
second batches.
11. What compound is the third batch of crystals? Dispose of the chemicals in the designated
receptacles.
Pre-Lab 4: Metathesis – To Exchange or Not
Hopefully here for the Pre-Lab Name(Print then sign): ___________________________________________________
Lab Day: ___________________Section: ________TA__________________________
This assignment must be completed individually and turned in to your TA at the beginning of lab.
You will not be allowed to begin the lab until you have completed this assignment.
1. Write molecular, complete ionic, and net ionic equations for the reactions that occur, if any, when solutions of the following substances are mixed: (a) nitric acid and barium carbonate
(b) zinc chloride and lead nitrate