For this experiment we studied an oxidation-reduction reaction of magnesium and hydrochloric acid solution. We compared the experimental measured amount of a product and the amount predicted by the theoretical calculation of a balanced equation: Mg (s) + 2HCl (aq) → MgCl2 (aq) + H2 (g) PROCEDURE First we obtained a strip of pre-cut magnesium ribbon, cleaned it with steel wool to remove any signs of oxidation from the strip (which would alter our results), and then weighed it (individual masses of the magnesium strips for each trial can be found in Table I below).
We filled a large beaker more than half way full with warm tap water. We allowed the water to sit for a while to bring it to room temperature. While the water equalized, we obtained a 100ml-graduated cylinder and filled it with tap water (also at room temperature), and created with a lid of plastic wrap. It was important when creating this lid that no air bubbles formed under the lid. Then we inverted the graduated cylinder in the beaker full of water and carefully removed the lid.
We folded up our magnesium ribbon and wrapped it with a short piece of copper wire so the magnesium would react and not break off during the reaction. We put this magnesium piece at the bottom of the beaker underneath the graduated cylinder. Using a thistle tube we gently propped the end of the tube up under the inverted graduated cylinder and then poured 10ml HCl into it. As the magnesium began to react, we slowly pulled the tube out and waited as the gas bubbles began to displace the water in the cylinder. After all the magnesium was consumed, we moved the graduated cylinder up and down in the water to equalize the water levels. Once we had accomplished this we took a reading of the amount of gas collected in the cylinder. We repeated this two more times.
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RESULTS TABLE I: GAS MEASUREMENT FOLLOWING REACTION TRIAL I TRIAL II TRIAL III Mg Strip Mass: 0.07g Mg Strip Mass: 0.07g Mg Strip Mass: 0.08g Volume of Gas 80ml Volume of Gas 82ml Volume of Gas 87ml H2O Temperature 24˚C H2O Temperature 24˚C H2O Temperature 24˚C Barometric Pressure 75.1cm Barometric Pressure 75.1cm Barometric Pressure 75.1cm TABLE II: CALCULATION OF NUMBER OF MOLES OF H2 GAS COLLECTED (PV = nRT) P = 0.959atm, V = 0.080L, 0.082L, 0.087L respectively, R= 0.08206 L atm K-1 mol-1, T = 273.15 + 24˚C = 297.15K n= 3.147 x 10-3 mol. H2 n= 3.226 x 10-3 mol. H2 n= 3.422 x 10-3 mol. H2 TABLE III: CALCULATION OF GRAMS AND MOLE OF MAGNESIUM CONSUMED DURING EACH TRIAL BALANCED EQUATION: Mg + 2HCl → MgCl2 + H2 TRIAL I (3.147 x 10-3 mol. H2 ( 1mol Mg / 1mol H2) (24.3050 g/mol-1/ 1mol Mg) = 0.077g Mg Consumed (4 significant figures)
Convert mass in grams to moles: 0.077g Mg (1mol. Mg / 24.3050 g/mol-1) = 3.168 x 10-3 mol. Mg (4 significant figures) TRIAL II (3.226 x 10-3 mol. H2 ( 1mol Mg / 1mol H2) (24.3050 g/mol-1/ 1mol Mg) = 0.079g Mg Consumed (4 significant figures) Convert mass in grams to moles: 0.079g Mg (1mol. Mg / 24.3050 g/mol-1) = 3.250 x 10-3 mol. Mg (4 significant figures) TRIAL III (3.422 x 10-3 mol. H2 ( 1mol Mg / 1mol H2) (24.3050 g/mol-1/ 1mol Mg) = 0.083g Mg Consumed (4 significant figures) Convert mass in grams to moles: 0.083g Mg (1mol. Mg / 24.3050 g/mol-1) = 3.415 x 10-3 mol. Mg (4 significant figures) DISCUSSION OF RESULTS The result of calculating the experimental stoichiometry in comparison to the theoretical reaction given by the balanced equation above appears there are fewer moles (e.g. 0.003415 mol. Mg in Trial III) than there existed in the equation (1mol.Mg).
The sources of uncertainty in our experiment leading to this error could be several different possibilities.
One possible error that we notice could be that fragments of our magnesium ribbon broke away from the rest of the reacting magnesium and floated to the top. Although it appeared these particles eventually reacted with the HCl, it is possible they didn’t, therefore giving us an inaccurate measurement. Another possible source of uncertainty in our experiment is that some bubbles of Hydrogen gas evaded the mouth of the graduated cylinder and were lost in the beaker. Although we were careful about removing the thistle tube from under the graduated cylinder once we added the HCl, there isn’t a seal to insure that no bubbles escape once the Magnesium starts reacting. Therefore, my conclusion is that the measurement of the gas was inaccurate in all our trials to the point that it consequently altered the correct calculation of the amount of magnesium consumed and moles consumed.
Experiment 3: Evaluation of Gas Constant Purpose: The purpose of this lab is to demonstrate the ideal gas law under ordinary conditions. In this lab, the variables in the ideal gas law are known or can be found aside from the constant R. Thus, the R values can be found and relatively determine the relevancy of the ideal gas law to the lab conditions. The ideal gas law was tested using the ...
Significant figures are part of increased accuracy for this experiment. The rule that is relevant when calculation in this experiment is the multiplication and division rule which states: When multiplying or dividing, the result must be reported with the same number of significant figures as the measurement with the fewest significant figures. Measurement, is the key word in this rule. One needs to account for the measured calculations when deciding how many significant figures are necessary, and then carry this through all the way to the end of your calculations. It is noted in the examples in the tables above how many significant figures in each calculation.