CHEMISTRY LAB REPORT

Lab 2: Hydrates
Chapter 3 of Chemistry, Principles and Reactions by Masterton and Hurley; 8th Edition contains most of the background needed for this lab. Read all of Chapter 3 prior to coming to the lab. Section 3-2 deals with percent composition calculations and section 3-3 has information of writing chemical equations that may be useful. On page 71 there is a box Chemistry Beyond the Classroom which gives additional information on hydrate chemistry.
OBJECTIVES
In this experiment you will apply some basic ideas of chemistry to the study of hydrated compounds. You will observe changes in the structure of some hydrates due to changes in the number of water molecules in their hydration shell. Simple chemical tests will be used to identify the component ions of an unknown hydrate.
THEORY
Most solid chemicals will contain some water if they have been exposed to the atmosphere. In some cases, the water present is absorbed only on the surface of the compound and is easily removed by gentle heating. In other cases, the water present is incorporated as an essential part of the crystal structure of the compound; it is generally bound to the cation. The water present in this second group of salts is called the “water of hydration” or “water of crystallisation.” It is usually present in stoichiometric amounts, for example, AlCl3∙6H2O(s). Some hydrated compounds spontaneously gain or lose water from the atmosphere. Such compounds are respectively described as deliquescent, or efflorescent.
Deliquescent compounds absorb water from the atmosphere so strongly that they might dissolve in their own water of hydration. They are said to be hygroscopic.
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Industrially, some of these compounds, such as CaCl2(s) or P2O5(s), are widely used to dry liquids or gases. As a result, they are also called desiccants. For these compounds, the vapour pressure of water exerted by the solid is less than that of the (partial) vapour pressure of water in the surrounding atmosphere. Alternately, if we know that a compound deliquesces, we can conclude that the vapour pressure of water exerted by the solid is less than that of the (partial) vapour pressure of water in the surrounding atmosphere. Efflorescent compounds lose water spontaneously if they are open to the atmosphere. For these compounds, the vapour pressure of water exerted by the solid is greater than that of the vapour pressure of water in the surrounding atmosphere. As before, if we know that a compound effloresces, we can conclude that the vapour pressure of water exerted by the solid is greater than that of the vapour pressure of water in the surrounding atmosphere.
Most hydrates are stable at room temperature. To remove the water of crystallisation from such a hydrate, it must be heated to a higher temperature where the vapour pressure of water around the solid hydrate exceeds the partial pressure of water in the atmosphere above it. At this higher temperature, the water of crystallisation is driven off and dehydration occurs. When all of the water has gone, the salt is said to be in its anhydrous form. The temperature required to dehydrate a salt varies considerably from hydrate to hydrate, depending on the vapour pressure of water around the hydrate. For example, Na2CO3∙10H2O(s) exerts a vapour pressure of 24.2 torr at 20°C. If the partial vapour pressure of water in the atmosphere is greater than 24.2 torr at 20°C, this compound will need to be heated until the vapour pressure of the compound exceeds the vapour pressure of the atmosphere in order to drive off its water of hydration. However, if the partial pressure of water in the atmosphere is less than 24.2 torr, Na2CO3∙10H2O(s) will lose water slowly; that is, it will effloresce. Since the partial pressure of water is 17.5 torr at 20°C, (see Chemical Data Sheet, Table 13), it will effloresce slowly.
As a hydrate is heated, and the water of hydration is driven off, the compound may undergo one or more of the following changes:
It may undergo a colour change. For example, CoCl2∙6H2O(s) is red,
CoCl2∙2H2O(s) is violet, and CoCl2(s) is blue. The change in colour is due to the
change in number of water molecules bound to the cation.
The form of the crystal may change. If it does, a dry powdery solid usually results.
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c. It may dissolve in its own water of crystallisation as the crystal structure breaks down. By continued heating, this water will evaporate, leaving the anhydrous solid behind.
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EXPERIMENTAL METHOD
In this experiment, you will watch a video of a few simple tests on hydrates to study their properties. You will then be given an unknown hydrate and photos of a series of tests that will allow you to determine which ions are present. A crucible and a portion of unknown sample will be weighed and then the water will be driven off by heating to obtain the anhydrous salt. By weighing the dried salt in the crucible, you will be able to determine how much water is lost and, when you have determined which ions are present, you will then be able to calculate the empirical formula of the hydrate (examples of such calculations are given in your textbook).
REAGENTS AND EQUIPMENT
Hydrate samples (unknowns) crucible, cover and tongs Lab balance (to 0.001 g) Test tubes
CuSO4∙5H2O(s) (crystals) Sugar (sucrose: C12H22O11(s))
EXPERIMENTAL PROCEDURE
Platinum wire 6M HCl(aq) 1M NaOH(aq) 1M BaCl2(aq) NaCl(s)
CdCl2∙3H2O(s)
Bunsen Burner 6M HNO3(aq) 1M H2SO4(aq) 0.1 M AgNO3(aq) NaBr(s)
Anhydrous CaCl2(s)
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Note: The procedure is presented as though you were performing the experiment in the lab. This is the same procedure that was followed in the video.
I. Unknown Hydrate
Have you remembered to put on your safety glasses? Set up as shown in
Figure 1, a Bunsen burner under an iron ring (the iron rings are from the
side storage cabinet). Put the wire gauze on the ring.
Obtain a hydrate sample from your Lab Instructor and record the code
number on the Observations Sheet. Describe the appearance of the hydrate on your Observations Sheet. Some of these samples decompose on heating, or give off a lot of water. You might want to check the behaviour of your
unknown by putting a little bit of it into one of your smallest test tubes, and
heating gently with your Bunsen burner.
(Note: After heating, crucible tongs should be used for moving the crucible and lid.)
Wire Gauge
Crucible Lid
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should be allowed to circulate)
Figure 1. Heating a Crucible (air
Wash your crucible and lid, wipe them dry, and put them side by side on the wire gauze. Heat the clean crucible and cover intensely for about 2 minutes to dry it well. (Note: Instructions for use of the Bunsen burner are given in the Appendix, page AP-4 of this manual. Read the instructions, and then ask your Lab Instructor for assistance if needed). Allow it to cool on your wire gauze. Do not place the crucible or cover on the bench top. During the cooling period, proceed to do Part II and III. When cooling is complete (crucible has reached room temperature), weigh the crucible and cover on the top-loading balance. Record the weight to 0.001 g. Remove the crucible and cover from the pan on the top-loading balance and add 3 – 4 g of the hydrate. Weigh the crucible and cover and the hydrate together, again to 0.001 g. If the weight of the contents is not in the 3 – 4 g range, remove the crucible from the balance again and use your spatula to add or remove compound. UNDER NO CIRCUMSTANCES ARE YOU TO TRANSFER CHEMICALS WHILE THE CRUCIBLE IS ON OR OVER THE BALANCE PAN. (In all of the experiments, you must
Crucible
Burnsen Burner
Tripond Stand
remove the containers from the balance pan of the top-loading balance before adding or removing chemicals). Re-weigh the crucible and cover and hydrate. Repeat the weighing procedure until the hydrate weight is in the desired range. Any excess hydrate is placed in the “Waste Solid” container by the balances.
C. Place the crucible (with the hydrate inside) on the wire gauze. Carefully put the lid beside the crucible. Heat the crucible and contents, gently at first, until you observe whether the salt melts and dissolves in its own water of crystallisation. If it does become a liquid, heat very gently around the edges at first to avoid splattering. Then, increase the heat moderately, but do not allow the crucible to become red hot. You will recognise when the hydrate is dry by noting (a) a colour change, (b) when the liquefied sample has completely dried, or (c) when no more changes are apparent in the dried powder. In some samples, excessive heating will decompose the compound. Such samples will be labelled. If this becomes noticeable, stop heating immediately.
D. Allow the crucible and contents to cool on your wire gauze to room temperature. Make sure to put the lid back on the crucible while it is cooling, using crucible tongs, re-weigh the crucible and its contents quickly to avoid re-absorption of water from the air by the salt. After obtaining the weight, reheat the crucible (take the lid off!) and contents a second time. This time, do not heat as strongly or as long as before. If decomposition occurs (marked colour change), stop immediately. Allow to cool, and record the weight again to determine if a constant weight has been reached. If constant weight is not achieved, heat your sample for a third time. When finished, put the solid in the solid waste container provided in the balance room by the electronic weighing scales.
II. Properties of Some Hydrates
A. Copper(II) sulfate pentahydrate
1. Heat a few crystals of copper(II) sulfate pentahydrate in a dry test
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tube held (with a test tube holder) in an inclined position over a Bunsen burner flame. Do not point an open test tube towards anyone as you gently move it back and forth through the flame. What evidence is there that water was present in the crystals?
Continue heating until no further change is observed. How does the residue differ from the original substance in appearance and composition?
Using your Pasteur pipette, add two or three drops of water to the residue. What do you observe? (Can you explain what you see?)
Pour the waste into the “Inorganic Aqueous Heavy Metals” (I.A.H.M.) container in the waste fume hood.
Sugar (sucrose: C12H22O11(s))
Note that the formula for sucrose is not (C12∙11H2O). Units of water molecules are not present in sucrose. (Look up the structure of sucrose in your textbook.)
Get a test tube from your Lab Instructor for this part of the
experiment. Put a few crystals of sugar into it. There will be a smell of caramel (similar to toasted marshmallows). Heat the sugar following the previous procedure (A.1 and A.2). Is there evidence that water is present?
Continue to heat until a brown/black solid residue is formed in the test tube (similar to burnt marshmallows). Using a Pasteur pipette add a few drops of water to the residue. What do you observe? Pour the liquid into the sink, and return the test tube to the tray.
CdCl2∙3H2O(s)
In the laboratory some CdCl2∙3H2O(s) in an open bottle has been exposed
to the air for several hours. What difference do you observe between the appearance of the crystals in the sealed and open bottles?
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D. Anhydrous CaCl2(s)
Observe the appearance of some anhydrous CaCl2(s) that is in a sealed bottle and some that has been exposed in a bottle open to the air for several hours. What difference do you observe?
III. Identification of ions in hydrate samples.
To identify your ions, dissolve about 0.5 g of your unknown hydrate in about 10 mL of water. Record the colour on your observation sheet. This is your unknown hydrate solution. Then carry out the following tests.
A. Anion analysis: CO32-, SO42-, Cl-, Br- (see page 2-15)
Proceed in the order outlined below. Once a positive anion test is obtained, you will not need to proceed any further with the anion analysis, as further testing will actually give wrong results.
WARNING: 6M HCl(aq) is corrosive. To a test tube containing 1
mL of unknown hydrate solution prepared above, add 1 mL of 6M HCl(aq). Note any gas being produced. The evolution of a gas indicates the presence of CO32-(aq). If gas is evolved, pour the waste into the I.A.H.M. container in the fume hood.
If no gas is evolved in III.A.1. add 1 mL of BaCl2(aq), one drop of 6M HCl(aq), and 2 mL of H2O(l) to the salt solution from III.A.1. A precipitate of BaSO4(s) confirms the presence of sulfate. Pour the waste into the I.A.H.M. container in the fume hood.
This step requires a fresh sample of the unknown hydrate solution. If there has not been a reaction in III.A.1. or III.A.2., add three drops of AgNO3(aq) and one drop of 6M HNO3(aq) (WARNING: 6M HNO3(aq) is corrosive) to 1 mL of your unknown hydrate solution to confirm Cl-(aq) or Br-(aq). A white precipitate confirms chloride, a cream precipitate confirms bromide. If you wish to check the colours of AgCl(s) and AgBr(s) run this test on 50 mg of NaCl(s) and 50 mg of NaBr(s) and compare the results with those of your
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unknown. If the colour of the silver precipitate is masked by the colour of the cation, decant the liquid. Then wash the precipitate with water and discard the wash liquid. Dispose of waste in the I.A.H.M. container.
B. Cation Analysis (see p. 2-15)
1. Compare the colour of the unknown hydrate solution with the
colours listed in the cation table.
2. WARNING: 1M NaOH(aq) is mildly corrosive. To 1 mL of the
unknown hydrate solution, add 3 drops of 1M NaOH(aq). Record your observations. Then continue to add 1M NaOH(aq), drop by drop, until about 2 mL have been added. Stir. Record your observations. Compare your observations with the possibilities given in the third (B.2.) and fourth (B.3.) columns of the cations chart on page A-15. Dispose of the waste in the I.A.H.M. container.
3. To distinguish barium from magnesium, a few drops of 1M H2SO4(aq) (WARNING: 1M H2SO4(aq) is mildly corrosive) can be added to 1 mL of your original salt solution. A white precipitate confirms barium since MgSO4(aq) is soluble but BaSO4(s) is not. Dispose of the waste in the I.A.H.M. container.
4. If no precipitate has formed after adding NaOH(aq) or H2SO4(aq), the ion must be Na+ or K+. To distinguish between Na+ or K+ ions, or to confirm one of the other ions, a flame test can be carried out. The wire must first be cleaned by alternately dipping a platinum wire into concentrated HCl(aq) and then heating it in the hottest part of the Bunsen burner flame until the wire gives little or no colour to the flame (platinum wire can be obtained from your Lab Instructor.) If you are not sure where the hottest part of the flame is, see the instructions in the Appendix, AP-4. Repeat this several times until the wire is clean. Place a solid crystal of the unknown on the clean wire loop, heat it in the Bunsen burner flame and observe any colour
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in the flame.
Before leaving the Lab, check to see that all your equipment is back in your locker, and that the lock is securely locked. When you close the lock, spin the dial on the lock; otherwise the lock might still open. You will have to pay for any missing items. If a Lab Instructor finds your locker unlocked, it will be put on backwards to let you know that it had been found opened.
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INFORMATION TO IDENTIFY ANIONS (PART III A.)
Anion
A.1.
add 1 mL of 6M HCl
A.2.
add 1 mL of 1M BaCl2 + 1 drop HCl
A.3.
add 6 drops of 1M AgNO3 in 6M HNO3
gas evolved
CO32- SO42-
Cl- Br-
–
– –
no reaction no reaction no reaction
white ppt.
no reaction no reaction
white curdy ppt.
cream-coloured ppt.
INFORMATION TO IDENTIFY CATIONS (PART III B.)
Cation
B.1. colour of solution
B.2.
add 3 drops of 1 M NaOH
B.3. add 2 mL of 1 M NaOH
B.4.
add 6 drops of 1 M H2SO4
B.5. Flame Test
blue
green
yellow
Mn2+
colourless or pale pink
white ppt. turns brown
ppt. stays
no reaction
–
violet
ppt. dissolves
no reaction
Co2+
pink
blue ppt., turns pink→brown
brown ppt. stays
brown ppt. stays
–
white ppt.
Cu2+
Ni2+ Na+
K+
Al3+ Mg2+
Ba2+ Note: ppt. is
blue ppt. green ppt.
no ppt.
no ppt. white ppt. trace of white ppt.
colourless
an abbreviation for precipitate
ppt. stays Ppt. stays
no ppt.
no ppt. ppt. stays
ppt. stays
no reaction no reaction
no reaction
no reaction no reaction
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