Download Reactions of common metals and properties of

Experiment 3
Metals and Hydrogen
Reactions of common metals and properties of hydrogen
"Since I ran down that coiner by the zinc and copper filings in the seam of his
cuff, they have begun to realise the importance of the microscope."
Sir Arthur Conan-Doyle, "THE ADVENTURE OF SHOSCOMBE OLD PLACE"
Background information
Group 1: alkali metals
The first group of the periodic table contains the elements lithium, sodium, potassium, rubidium,
and cesium, and the short-lived radioactive element, francium. They are called the alkali metals
because their hydroxides are all soluble bases or alkalis. The most important common
characteristics of these elements are the single electron in the outermost energy level of their
atoms, and the comparatively low energy required to detach this electron from the atom, as
shown by the ionization energies.
Element
H
Li
Na
K
Rb
Cs
Valence e- configuration
1s1
2s1
3s1
4s1
5s1
6s1
First ionization energy (kJ mol-1)
1312
520
496
419
403
376
Hydrogen is clearly the exception in this series. Accordingly, hydrogen is a diatomic gas,
whereas all the others are metallic solids. In some ways the chemistry of the alkali metal cations
resembles that of the proton, H+, but there are more differences than similarities. Hydrogen also
forms an anion, the hydride ion, H-, and many metals, including the alkali and alkaline earth
metals form salt-like hydrides, such as NaH. Such hydrides have similar crystal structures to
alkali halides. However, they are extremely reactive towards water, forming the corresponding
metal hydroxide and liberating hydrogen gas. These salt-like hydrides find some use as reducing
agents in organic chemistry.
Although the outer electron is easily removed for the alkali metals, the underlying ones can be
removed only by very high energies (why?); hence ordinary chemical reactions involve the loss
or transfer of only a single electron per atom, as illustrated by the equation: M → M+ + e-.
Accordingly these elements invariably display an oxidation state of +1 in their compounds.
Many of the physical and chemical properties of these elements may be correlated with the
structures of their atoms. Thus in the solid state the electrons are held so loosely that they can
readily pass from association with one atom to another under the impulse of an electric potential;
hence these elements show in a high degree the conductivity and other properties characteristic
of metals.
Since the valence electrons are held so loosely, it is not surprising that the alkali metals are the
most active of all metals. They tarnish in air because the metals lose their valence electrons to
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Experiment 3
Metals and Hydrogen
oxygen and they react violently with water. Sodium, for example, transfers electrons to
hydrogen ions in water, thus liberating hydrogen gas.
2 Na Æ 2 Na+ + 2 e2 H+ + 2 e- Æ H2(g)
The two half-reactions combined can be written as:
2 Na + 2 H+ Æ 2 Na+ + H2(g)
Atoms of the alkali metals are easily excited; even the flame of a Bunsen burner can excite their
valence electrons. As the electrons jump back to lower energy levels, they give characteristic
colours to the flame; lithium imparts a red colour, sodium a yellow colour, and potassium a lilac
colour. The colours are so pronounced that even if only traces of an element are present, the
metals can readily be recognized by a flame test.
The sodium flame test is capable of detecting 10-10 g of sodium. Since this quantity of sodium is
present in practically every substance, it is necessary in making a flame test for sodium to judge
the amount of sodium present largely by the length of time that the intense yellow colour persists
in comparison to pure sodium chloride as a standard. The potassium flame test is about five
thousand times less sensitive than the sodium test. In the presence of any considerable amount
of sodium it is necessary to use a thick blue glass to cut out the yellow sodium light in order to
detect the potassium.
Group 2 (alkaline earth) and group 12 metals
The elements with two valence electrons may be divided into a main group consisting of
beryllium (formerly often called glucinium from the sweet taste of some of its compounds),
magnesium, calcium, strontium, barium, and radium, and a subgroup (group 12), zinc, cadmium,
and mercury. As a common characteristic the group 2 elements have core electrons with the
same number and arrangement of electrons as the preceding noble gas, while the group 12
elements have cores with eighteen electrons in the outermost core shell.
Element
Be
Mg
Ca
Sr
Ba
Zn
Cd
Hg
Valence e- configuration
2s2
3s2
4s2
5s2
6s2
4s23d10
5s23d10
6s23d10
First ionization energy (kJ mol-1)
899
738
590
549
503
906
868
1007
Beryllium and magnesium are smaller and have much more acidic cations than the rest of group
2, and before the fundamental differences in atomic structure were recognized, they were
frequently classified with the group 12 elements. The remaining elements of group 2 have
hydroxides with distinctly alkaline properties and have long been known as the alkaline earth
elements, and we shall use the term as applying to all of group 2.
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Experiment 3
Metals and Hydrogen
In all of their compounds the elements of group 2 are present in the +2 oxidation state, i.e. both
of the two outer electrons are always lost when the elements enter into chemical reactions. It
will be observed (see tables above) that the ionization energy of the heavier members of the
group is as low as that of the alkali metals. Thus barium readily reacts with water, losing its
electrons to the hydrogen of the water and forming barium ion and molecular hydrogen:
Ba + 2 H2O Æ Ba2+ + 2 OH- + H2
The mechanism of the reaction involves taking the two electrons away from the barium and this
requires about 500 kJ mol-1 for the first electron, and 1000 kJ mol-1 for the second. The rate of
reaction of the group 2 elements with water is much slower than that of the alkali metals, even in
the case of the more electropositive members; and magnesium and beryllium are scarcely
attacked by water alone at ordinary temperatures. At low temperatures the metal surfaces are
protected from rapid oxidation by oxide films; this is particularly true of beryllium and
magnesium. Beryllium does not dissolve with appreciable speed in water even when boiled.
Magnesium evolves hydrogen very slowly with cold water, while with calcium a slow stream of
bubbles is observed. The metals all dissolve rapidly in acids.
Beryllium and magnesium compounds do not ionize sufficiently to impart colour to the Bunsen
flame. Volatile calcium compounds give a brick red, strontium a carmine, barium a yellow
green, and radium a crimson flame. The colouration is very intense with the chlorides or
nitrates, but is not satisfactory with the oxides or sulfates due to their low volatility.
Water containing soluble calcium and magnesium salts is known as hard water. It is
objectionable in the laundry because soap, which contains sodium or potassium salts of the
higher fatty acids, such as stearic or palmitic acid (C17H35COOH and C15H31COOH), forms
insoluble salts with calcium and magnesium. This results in a waste of soap and the precipitate is
a slimy curd that is difficult to remove. In regions where these salts in the water are largely
bicarbonates, softening occurs by boiling, which precipitates CaCO3 (boiler scale). In regions
where they are sulfates or other salts, they can only be removed by distillation or (more
economically) by ion-exchange. In this process (i.e. inside a water softener), calcium and
magnesium ions are left on the ion exchanger, and sodium ions (from cheaply available rock
salt) are put into solution. Sodium stearate and palmitate are soluble, and hence do not give
slimy precipitates.
Zinc and mercury are members of group 12. The much higher first ionization potential of zinc
and mercury are reflected in the lack of reactivity of these metals with water. They do, however,
dissolve in dilute mineral acids, with the evolution of hydrogen. Zinc has one positive oxidation
state, +2. The metal is often used as a strong reducing agent; it is readily oxidized, with
formation of hydrogen gas, in both acidic and basic solutions:
Zn + 2 H+ Æ Zn2+ + H2
Zn + 2 H2O + 2 OH- Æ Zn(OH)42- + H2
This amphoteric nature of zinc and its salts is also displayed by aluminum, a close relative in the
periodic table.
Mercury has two positive oxidation states, +1 and +2. In noncomplexing aqueous solution,
mercury(I) exists as the dimeric mercurous cation, Hg22+, and mercury(II) exists as the mercuric
cation, Hg2+. The Hg2+ ion reacts with metallic mercury:
Hg2+ + Hg → Hg22+
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Experiment 3
Metals and Hydrogen
The equilibrium constant for this reaction (that is, the equilibrium ratio of [Hg22+]/[Hg2+]) is 83.
The reaction may be shifted by the addition of a reagent that reacts preferentially with one of the
ions to form a stable complex or a slightly soluble compound (i.e. a precipitate).
Pre-lab preparation
Read carefully through the background information and the indicated chapters in the references
at the end. You should know the material contained in any general chemistry texts (e.g. Kotz
and Treichel), and you can test yourself by doing the problems.
SAFETY
1.
Wear approved eye protection at all times.
2.
Use caution at all times during this laboratory. Some dangerous chemicals will be used!
3.
Sodium, and especially potassium, are extremely dangerous when in contact with moisture.
Never use more than the suggested quantity of metal in the experiments. Cut the metals
while still under the parafin. Note that potassium is spontaneously flammable in air!
4.
5.
All procedures except #2 must be done in a fume hood. Follow all instructions.
Be courteous in the use of stock reagents. Replace containers promptly to the side
counters so that others can find them. Follow the instructors guidance in staggering your
progress through the experiments.
Procedure
1.
Reactions of active metals with water
Wear safety glasses! In the FUME HOOD, add a piece of sodium smaller than a pea
to a 250 ml beaker filled to about three-fourths capacity with water. ENSURE THAT A
SAFETYSHIELD IS BETWEEN YOU AND THE BEAKER. USE A LONGHANDLEDSPATULA OR TONGS TO PLACE THE METAL IN THE BEAKER.
MAKESURE THE REACTION IS COMPLETLY OVER (ALL METAL
CONSUMED)BEFORE CONTINUING. When the reaction stops add several drops of
phenolphthalein indicator which turns pink in the presence of a base and is clear in acid
solution.
Repeat the above procedure with a piece of potassium metal no larger than a pea!.
Repeat the above procedure with calcium metal. It may be necessary to scrape the oxide
layer off the calcium first.
Cut a 3 cm strip of magnesium. Scrape off the coating of oxide with sandpaper or a file
until the strip is shiny. Add the metal to 25 ml of water in a small beaker. Heat the water
to boiling and then add several drops of phenolphthalein.
Repeat the same procedure with a piece of zinc metal. Is there any reaction? Did you
expect one?
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Experiment 3
Metals and Hydrogen
Write equations for the observed reactions and form a relative reactivity series.
2.
Flame tests for metal ions
This procedure can be carried out on the open bench. This will make it a lot easier to
obtain a steady Bunsen flame. Be careful with the flame and any flammable materials or
reagents!
Heat a platinum wire in a Bunsen flame and, if it causes a change in the colour of the
flame, clean the wire. The wire may be cleaned by dipping it in concentrated HCl in a
test tube and then heating it to redness until no colour change occurs in the flame. Save
the HCl for later cleaning.
Dip the wire into a few drops of 0.2 M Ca(NO3)2 solution in a test tube and then heat it in
the flame. Note the colour imparted to the flame by the calcium. Clean the wire.
Repeat the flame test successively with 0.2 M solutions of Sr(NO3)2, Ba(NO3)2, LiCl,
NaCl and KCl.
Repeat the sodium and potassium flame tests looking through didydium glass. Test the
mixed solution of Na+ and K+with and without the didymium glass.
3.
Reactions of metals with acid and alkali; identification of hydrogen
In the FUME HOOD, add a spatula-tip full of granular (a) zinc and (b) aluminum metal
to a 6.0 M HCl solution in a large test tube. Agitate the test tube, while keeping it
pointed away from you. Observe what happens when reaction starts.
Using a Bunsen burner, light a 12 cm wooden splint, using tongs to hold it. Hold the
burning splint in front of the mouth of the test tube while reaction with the metal
continues.
In the FUME HOOD, add a spatula-tip full of granular (a) zinc and (b) aluminum metal
to a 10.0 M NaOH solution in a large test tube. Warm the test tube in a Bunsen flame,
while agitating it and keeping it pointed away from you. Stop heating once the reaction
is initiated. Don't be fooled bysteam coming off! Observe what happens when reaction
starts.
Light a 12 cm wooden splint, and hold the burning splint in front of the mouth of the test
tube while reaction with the metal continues.
The presence of hydrogen in any of these reactions should result in an audible "pop"
when ignited. Do not use large pieces of metal to avoid dangerous quantities of
hydrogen!
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Experiment 3
4.
Metals and Hydrogen
Reaction of iron with oxidizing agents
In the FUME HOOD, light a Bunsen burner. Sprinkle a small quantity of iron filings in
the flame. What do you observe? Hint: this is not a flame test for iron, i.e. the emission is
not due to atomic transitions in the iron.
In the FUME HOOD, set up a small crucible above a Bunsen burner, using a crucible
stand. Put in about 0.25 g of iron filings, and about 1.0 g of powdered sulfur. Mix the
solids. Heat the mixture with the burner. Observe. Discontinue heating when emissions
from the crucible stop. What is the product?
5.
Reactions of the copper(I) and copper(II)
Dissolve a spatula-tip full of CuSO4 in 5 mL of distilled water in a test-tube.
What do you observe? What is the colour due to? Add a few drops of 6 M NaOH to the
solution. Observe what happens. Identify the products.
Put several spatula-tips full of Cu2O in a large test-tube, and add 10 mL of distilled water.
What do you observe?
Now heat the large test tube in a strong Bunsen flame until there is definite evidence of a
change occurring. Record your observations. If no change occurs add a couple of drops
of c.H2SO4. Continue heating until a change is observed.
Decant off the liquid into a clean test tube. What does this solution resemble? Add a few
drops of 6 M NaOH to the solution, and observe what happens.
Carefully look at the solids left in the heated test-tube. What do they resemble?
Report
In this and the subsequent descriptive chemistry labs, you must record detailed observations in
your lab notebook (see the Introduction to the lab manual for more details.) You should always
attempt to write a balanced chemical equation for each procedure that you do. One of the main
advantages of doing so, and of making it a habit to do so, is that the act of balancing often helps
you to discover "missing ingredients", i.e. species involved in the reaction which you have
overlooked. These notebooks will be evaluated on an on-going basis as well as at the end of the
semester.
The following are some of the observations you should be looking for:
• is a gas evolved (evidenced by bubbles rising in the liquid) other than that caused by boiling?
• is a vapour evolved (evidenced by a mist forming above the test-tube, as the vapour
condenses in the cooler air)?
• is a cloud evolved (i.e. an aerosol of liquid or solid particles ejected into the air above the testtube, evidenced by its opaque nature)?
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Experiment 3
Metals and Hydrogen
• is the material ionic or molecular (evidenced by the kind of solvent it dissolves in,
remembering that polar, protic solvents often dissolve ionic salts, while non-polar liquids
such as benzene or CCl4 dissolve only neutral molecules)?
• has there been a pH change (litmus paper or indicator paper or solution)?
• is heat evolved or absorbed? Note that this observation can only be made if no external
heating or cooling is applied. It must be a spontaneous change.
• has there been a colour change?
• has a precipitate formed, or has a solid dissolved? For this purpose, precipitate is any process
that causes a homogeneous solution to become non-homogeneous; some solids are
precipitated in extremely fine particles that do not settle easily. Another possibility is that a
colloidal suspension has formed. Use rigorous logic in making your judgments, avoiding
half-baked observations that you cannot be sure off afterwards.
The following are some of the principles governing inorganic reactions, which might help you to
explain your observations:
• has an acid-base reaction occurred?
• is there a possibility of an oxidation/reduction reaction (extremely common in these
experiments)?
• or some combination of the above? Consult the Latimer diagrams for help.
• what are the electronegativities of the reagents?
• could a Lewsi acid/base reaction have occured?
• could the observed changes be interpreted in terms of HSAB theory?
• can LeChatelier's principle explain the course of the reaction?
The exact form of the report you will be asked to provide will be specified by the lab instructor.
References
1.
D.F Shriver, P.W. Atkins and C.H. Langford "Inorganic Chemistry", NY: Freeman, 1990.
Ch. 9.
2.
R. Chang, "Chemistry, 3rd Edition", NY: Random House, 1988. Ch. 20, pp. 801-845; Ch.
21, pp. 849-854.
3.
P.W. Atkins "General Chemistry", NY: Scientific American Books, Ch. 18, pp. 660-689;
Ch. 21, pp. 771-779.
4.
J.C. Kotz and P. Treichel, Jr., "Chemistry and chemical reactivity", 4th edition, Saunders,
1999, Ch. 22, p. 1004 - 1021.
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