In the Laboratory
Determination of the Relative Atomic Masses of Metals by Liberation of Molecular Hydrogen W. Earle Waghorne and Andrew J. Rous* School of Chemistry and Chemical Biology, University College, Belfield, Dublin 4, Ireland; *
[email protected] There are a number of popular undergraduate experiments that involve the generation of gases from solids. In some cases, the experiment is qualitative. In others, an attempt is made to determine a relative atomic mass by relating the volume of the gas to either the mass of the gas (1) or the mass of a solid precursor (2–4). Experience shows that students are able to relate the obvious gas production to the mass change, making the chemistry both more approachable and more interesting. In this experiment, we measure the atomic masses of metals. This is the third experiment carried out by our first-year students. In the present experiment the relative atomic masses of the reactive metals, Mg, Ca, and Al, are determined by reacting weighed samples of each with HCl (4 M) and collecting the H2 gas evolved. The amount of hydrogen produced is calculated from the volume of hydrogen produced. The ratio of moles of metal reacted to moles of hydrogen produced can be determined from the balanced equation for the chemical reaction and so the number of moles of metal consumed can be calculated. The relative atomic mass is then calculable from the mass and number of moles of metal in the sample. Our method is to have students carry out the reaction with 0.5 g samples of each of the metals. This results in volumes of hydrogen gas ranging from 250 mL to 750 mL, providing a clear visual example that mass is not the relevant parameter for determining the amount of reaction that occurs. In our laboratories the experiment is carried out by two groups of students: one comprising students who have studied chemistry at secondary school and the second group comprising students who have not. While the fundamental experiment remains unchanged for these two groups, treatment of the information is slightly different. An example of this is in the conversion of the hydrogen volume to the number of moles. Students with secondary school chemistry are given the atmospheric pressure, provided by the local weather service, and carry out the calculation using this, the measured volumes, and water bath temperature via the ideal gas law. This provides a useful review of ideal gas behavior. In contrast, students who have not studied chemistry at secondary school are shown the ideal gas law (which is taught formally later in the course) and are given the volume occupied by one mole of gas at the temperature and pressure of the experiment. The pronounced exothermicity of the reactions provides a simple example of enthalpy changes in chemical reactions. Again this provides a point of departure for review with students who have studied chemistry at secondary school (where thermochemistry has been covered) or an introduction to the ideas for those without the relevant background. While the experiment is straightforward it provides a number of insights that instructors can usefully exploit. Thus, for example, there is a significant initiation time for the reaction of Al and at the end of the experiment the acid solution contains a gray precipitate. This observation illustrates the role of the oxide coating in protecting Al, which is an extremely reactive metal, and yet is widely used, for example in cookware. In contrast, Ca 350
and Mg, which leave no precipitate, begin to react immediately on contact with the acid but their reaction rates, as measured by the time it takes for bubbling to cease, are clearly different, showing that the reaction rate is not the same in all cases. In our experiments the difference in reaction rates reflects both the electronegativity and the state of the sample provided, Mg being in the form of ribbon while Ca is granular and so has a higher surface-to-volume ratio. The experiment also provides an informal way of demonstrating the behavior of gases, Thus raising or lowering the inverted collection vessel results in increases or decreases in the measured volume of gas, reflecting the changes in pressure owing to the differences in the water levels inside and outside of the collection vessel (Boyle’s law) while there is a small but observable decrease in gas volume when the flask is cooled at the end of the reaction (Charles’s law). Apparatus The apparatus consists of a 250 mL Büchner flask and rubber stopper with the outlet from the side arm directed, via a delivery tube, into a plastic measuring cylinder1 inverted in a water bath. A schematic of the setup in available in the online material. Method The students are provided with approximately 0.5 g samples of the three metals in closed sample bottles. The students determine the mass of the sample by difference, weighing the sample and sample bottle using a precision balance (±1 mg precision) and then reweighing the empty sample bottle after the sample has been added to the acid. The apparatus is set up with the graduated cylinder and the submerged part of the delivery tube filled with water. Hydrochloric acid, 25 mL of 4.0 M HCl, is added to the Büchner flask. The sample of metal is then added directly from the sample bottle and the stopper inserted into the flask as quickly as possible to prevent loss of hydrogen gas. This method of delivering the sample avoids having students handle the metal samples. Our experience has shown that Ca granules, Mg ribbon, and Al foil have suitable reactivities. The reaction is allowed to run until the evolution of gas has ceased, typically around 2 to 5 minutes. The Büchner flask is then cooled to the bath temperature by holding it in the water bath.2 Finally the graduated cylinder is raised or lowered until the liquid levels inside and outside the cylinder are equal and the volume of gas generated is read. The apparatus is then cleaned and reassembled and the experiment repeated for the remaining metal samples. There is a small but measurable volume of water displaced from the delivery tube during the experiment. This volume is estimated by emptying the volume of water in the delivery tube
Journal of Chemical Education • Vol. 86 No. 3 March 2009 • www.JCE.DivCHED.org • © Division of Chemical Education
In the Laboratory Table 1. Student Values for Relative Atomic Masses of Metals Groupa
Mg/(g/mol)
Ca/(g/mol)
Al/(g/mol)
1
23.8
27.5
25.7
2
26.2
42.7
25.9
3
27.3
29.6
25.0
4
23.1
41.5
24.1
5
26.0
40.8
23.1
6
27.8
43.5
26.3
7
23.0
38.3
25.7
8
22.3
40.3
25.3
9
26.4
41.0
25.9
10
22.7
43.4
26.7
11
26.5
41.6
26.5
12
24.3
40.5
19.8
13
24.1
41.5
27.4
14
24.7
37.2
20.2
15
24.5
40.1
26.8
16
26.5
36.7
27.5
17
27.6
41.8
26.2
18
27.8
34.5
24.1
19
23.5
37.7
25.0
20
24.5
39.6
25.5
21
27.0
40.1
27.5
22
25.2
40.9
26.7
23
27.6
34.5
25.6
24
23.2
40.6
24.1
24.5
36.7
27.2
Averageb
25
25.2 ± 1.8
38.9 ± 4.0
25.4 ± 2.0
Literature
24.3
40.1
27.0
aEach
set for the same student, chosen randomly except that 5 data sets containing obvious calculation errors have been excluded. bPrecision quoted as standard deviation.
at the initial setup into a 50 mL measuring cylinder. This volume is added to that measured previously. It is clear that the experiment can be extended to other metals by varying the concentration of the acid used. However, this introduces the possibility of reacting a metal sample with an overly concentrated acid solution and so we have restricted the experiment to metals that react at reasonable rates with 4 M HCl. This might be overcome by effecting the metal–acid reaction electrochemically (4).
Hazards Hydrochloric acid is corrosive and causes severe irritation and burns. Hydrogen gas is flammable and explosive. Metal samples react exothermically with water producing flammable and explosive gas. Results The results obtained for the relative atomic masses of the three metals by a random sample of 25 students are presented in Table 1. The precision of the mean value is shown as a standard deviation. The results in Table 1 are satisfactory given the experience of the students and the simplicity of the apparatus. The two principal sources of error are loss of gas, which leads to large values, and inappropriate rounding of masses and calculated values. Notes 1. We find that a 1000 mL graduated cylinder shortened to either 600 or 800 mL is suitable as a collection vessel. It is possible to work with smaller masses of metal and thus smaller volumes of gas but the experimental errors are proportionally greater. 2. The water bath is typically a 54 L plastic storage box, which is able to accommodate 2 sets of apparatus.
Literature Cited 1. Bunce, D. M.; Schwartz, A. T.; Silberman, R. G.; Stanitski, C. L.; Stratton, W. J.; Zipp, A. P. Chemistry in Context: Applying Chemistry to Society, 2nd Ed.; American Chemical Society: Washington DC: 1997; Chapter 6. 2. Peller, J. R. Exploring Chemistry: Laboratory Experiments in General, Organic and Physical Chemistry; Prentice Hall: Upper Saddle River, NJ, 1998; Chapter 10. 3. Beran, J. A. Laboratory Manual for Principles of General Chemistry, 5th ed.; Wiley: New York, 1994; Chapter 18. 4. Kildahl, N.; Varco-Shea, T. Explorations in Chemistry, A Manual for Discovery; Wiley: New York, 1996; p 63.
Supporting JCE Online Material
http://www.jce.divched.org/Journal/Issues/2009/Mar/abs350.html Abstract and keywords Full text (PDF) Supplement
Student handouts
Instructor notes including the hazard assessment form
© Division of Chemical Education • www.JCE.DivCHED.org • Vol. 86 No. 3 March 2009 • Journal of Chemical Education
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