Download Experiment 3: Some Reactions of the First Transition

Experiment 3
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Experiment 3:
Winter 2013
Some Reactions of the First Transition Series
Introduction
The transition elements are characterized by the formation of coloured complexes,
some of which you will see in other experiments. In this series of tests, you will see how the
ions of the most common oxidation states of certain first row transition elements behave in various standard conditions and also learn one or two unique reactions of each ion.
Test 1–2 mL samples of the solutions provided, unless otherwise stated. Sulfates,
nitrates and chlorides of these metals are usually quite soluble in water; the sulfate ion is pre ferred because it does not readily interfere with redox reactions.
Record all observations in the provided tables. Several of these tests are multistage.
Start each numbered test with a fresh portion of stock solution.
WASTE COLLECTION BOX:
Fluoride, metal ion solutions.
Procedure
a.
Reactions of titanium
The most stable oxidation state of this element is +4. Solutions do not contain the highly
polarizing Ti4+ ion, but partly hydrolyzed ions like Ti(OH) 2+2. The titanium compound used for
this test is dissolved in concentrated HCl and is stored in a fume hood. As this solution is
strongly acidic, HANDLE IT WITH CARE!!!! Use 2–3 mL of titanium tetrachloride for each test
listed below.
Ti 1.
Add 6 mol.L-1 sodium hydroxide dropwise until the volume is at least doubled.
Ti 2.
Add a small piece of zinc or tin metal to the acidic solution.
Ti 3.
Add 6 mol.L-1 sodium hydroxide solution dropwise with through mixing until a
trace of precipitate forms and then add a few drops of 6% hydrogen peroxide.
Pour part of this mixture into a second test tube and set it aside. Add a spatula end of sodium
or ammonium fluoride to one portion of the mixture. Stopper the test tube and shake to dissolve
the solid. Compare this test tube with the other portion of mixture set aside.
(CAUTION: Use gloves when handling fluorides with care!!!)
NOTE:
Pour the mixtures from the three tests above into the metal solution collection
box once the observations have been made.
NO equations are required for these titanium tests!!
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Procedure (continued)
b.
Reactions of chromium
Chromium (III) is the most stable oxidation state of this element in aqueous solution, but
the +6 state is also found in the well known chromate oxidizing agents. Use a 2–3 mL of chromium (III) chloride, nitrate or sulfate solution for each test below.
Cr 1. (a)
Add 6 mol.L-1 sodium hydroxide dropwise to 1 mL of the chromium solution
until excess is present (~ 2 mL).
Cr 1. (b)
Add 1 mL of 6% hydrogen peroxide to the basic solution from (a), place
the test tube in a beaker of hot water and wait until the solution is yellow.
Then add 1 mL of 1-pentanol (amyl alcohol), which will form a separate layer, 5 more
drops of hydrogen peroxide and acidify with 6 mol.L-1 sulfuric acid. Mix well. Note all
colour changes.
Cr 2.
Place 1 mL of fresh CrCl3 (aq) in a large test tube by adding a scoopula tip of
chromium(III) chloride to 1 mL of deionized water. Then add 2 mL of (6
mol.L-1 HCl) and a small spatula end of powdered zinc to this test tube. Next fit the test
tube with a Bunsen valve to keep out air. Leave the test tube to stand for about 5
minutes. Remove the valve and pour approximately half the solution through a cotton
plug and a glass funnel. Replace the valve in the original test tube. Compare the colours of the solution before (in the original test tube) and after exposure to air (the filtered
solution). Note all colour changes.
NOTE:
Dispose of the contents of the test tubes from these tests in a metal ion waste
collection box.
c.
Reactions of manganese
This is the first member of the series with a +2 cation which is stable under normal conditions. Use 2–3 mL of a manganese (II) sulfate solution for each test.
Mn 1.
Add 6 mol.L-1 sodium hydroxide dropwise until excess is present (about 2
mL). Leave to stand and note any subsequent colour changes.
Mn 2.
Acidify with an equal volume of 6 mol.L-1 nitric acid and add a tiny amount of
solid sodium bismuthate. If you see undissolved solid in the test tube, you have
used too much! Shake, leave to stand, and observe the colour of the solution.
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Procedure (continued)
c.
Mn 3. a.
Reactions of manganese (continued)
Dissolve two or three crystals of KMnO4 in 2 mL of 6 mol.L-1 NaOH. (NOTE:
The desired results will not be seen if too much KMnO4 is used). Let the test
tube sit for 1–2 minutes after mixing. Place a drop of this solution onto a clean
piece of paper towel and note the colour as the solution soaks into the paper
(NOTE: the paper towel is the reducing agent here. In this case including only
the reduction half-reaction in your lab report will be sufficient.). Observe for a
further 30 seconds and note any other changes.
Mn 3. b.
Add 1–2 mL of starch solution to the basic KMnO4 solution prepared above
and shake the test tube until a noticeable colour change occurs (NOTE: The
starch serves the same purpose as the paper towel did above.). If you do not
observe a noticeable change, add more starch solution to your test tube. Now
carefully acidify this solution with 3 mol.L-1 H2SO4. Filter the mixture by gravity
filtration and immediately observe the solution's colour and note the presence or
absence of a precipitate. If a precipitate forms, note its appearance.
d.
Reactions of iron
There are two oxidation states of iron which are reasonably stable under normal conditions. Iron (III) [ferric] is mildly oxidizing and iron (II) [ferrous] mildly reducing. Unless special
care is taken, solutions of either may contain traces of the other. A solution of iron (II) is not
provided for this reason; make up your own using iron(II) (ferrous) ammonium sulfate,
which seems to be more resistant to oxidation than the simple sulfate. For iron (III), a solution
of iron (III) chloride or nitrate can be used. Compare 2 mL portions of your solutions of each of
the two ions in the following series of reactions.
Fe 1.
To each solution add 6 mol.L-1 sodium hydroxide until excess is present.
Leave to stand for a few minutes to see if any further changes take place on
exposure to the air.
Fe 2.
(a)
Add 1 mL of potassium hexacyanoferrate (II) (aq) to solutions of each ion.
(b)
Add 1 mL of potassium hexacyanoferrate(III) (aq) to solutions of each ion.
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Procedure (continued)
d.
Fe 3.
Reactions of iron (continued)
Add a small crystal of potassium or ammonium thiocyanate to a portion of each
solution and shake to dissolve the solid. Pour a portion of the mixture into a
second test tube for each cation. Add a little solid sodium or ammonium fluoride to one
portion of each mixture. Stopper each test tube and shake to dissolve the solid.
Compare this test tube with the other portion of mixture set aside. (CAUTION: Handle
solid fluorides with care.) This is a sensitive test for iron (III).
NOTE:
Dispose of the mixture from Fe 3. in the fluoride collection bottle provided.
e.
Reactions of cobalt
When cobalt is complexed by nitrogen ligands, the +3 state is most stable but it is
strongly oxidizing when the ligands are water molecules. The familiar pink colour of cobalt ions
is due to the Co2+ (aq) ion. Use 2–3 mL of a cobalt(II) nitrate or sulfate solution for each test below.
Co 1.
Add 6 mol.L-1 aqueous ammonia dropwise until excess is present. Then add a
few drops of 6% hydrogen peroxide and leave to stand for a few minutes.
Co 2.
Add an equal volume of concentrated hydrochloric acid. Heat the acidified solution in a boiling water bath.
Co 3.
Add a spatula end of a solid thiocyanate, 2 drops of concentrated hydrochloric
acid and 1 mL of 1-pentanol. Shake to dissolve the solid.
f.
Reactions of nickel
Oxidation states of nickel other than +2 are quite difficult to obtain although a nickel (III)
oxide is used in rechargeable Ni / Cd batteries. Use 2–3 mL of a solution of nickel chloride for
each test below.
Ni 1.
Add 6 mol.L-1 aqueous ammonia dropwise until no further change is observed.
Ni 2.
Add a spatula end of a solid thiocyanate, 2 drops of concentrated hydrochloric
acid and 1 mL of 1-pentanol. Shake the test tube to dissolve the solid. Compare the results for this test with the results for test Co 3. above.
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Procedure (continued)
f.
Ni 3.
Reactions of nickel (continued)
Add 1 mL of dimethylglyoxime reagent and shake well. If no change is observed, then add 6 mol.L-1 aqueous ammonia dropwise until a change does occur (As well as being a diagnostic test for nickel, this precipitate can be used
quantitatively in the gravimetric analysis of nickel.).
g.
Reactions of copper
Copper is the only element in the series which has a stable +1 state, but even this disproportionates in aqueous solution and is only known in complexes or insoluble salts. A few,
highly oxidizing solid complexes of copper(III) are also known but by far the majority of compounds of this element contain copper(II). Use 2–3 mL of a copper(II) sulfate solution for each
test below unless a different volume is stated in the test procedure.
Cu 1.
Add 6 mol.L-1 aqueous ammonia dropwise with shaking until no further
change is observed.
Cu 2.
Add an equal volume of concentrated hydrochloric acid, then dilute to about
10 mL with deionized water and add a small piece of zinc.
Cu 3.
Prepare some saturated potassium iodide solution by dissolving ~ 3 g of KI in
2 mL of deionized water at room temperature. To 1 mL of 0.1 mol.L-1 CuSO4
(aq) in a test tube, add 3 or 4 drops of the saturated KI solution and then add dilute sodium thiosulfate dropwise to just remove the colour of the iodine. Starch
solution can be added to determine when the I2 has been removed. Next add ~
2 mL of saturated KI solution to the mixture in the test tube and shake. Note
what happens. More saturated KI solution can be added if necessary.
References
P. Atkins, T. Overton et al., Shriver & Atrkins Inorganic Chemistry, any edition.
G.I. Brown. Introduction to Inorganic Chemistry, second edition.
C. E. Housecroft and A.G. Sharpe. Inorganic Chemistry, any edition.
G. Svehla. Vogel's Qualitative Inorganic Chemistry, sixth edition.