Today's second lab was very exciting. Ameera and I got to experiment with actual chemicals, and act like real chemists! Our interesting tools were the best part.
The purpose of today's lab was to keep heating the hydrate CuSo4 .xH2O until much of the water evaporated, and the compound changed color from blue to a sky-tinted white. We needed to weigh the compound many times during the heating process to determine how the mass changed once more and more water evaporated. Then we needed to use our data once the water finished evaporating to determine the empirical formula the hydrate.
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| Before Heating |
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| After Heating (on the hot plate) |
Calculations:
1. Mass of the Hydrate Used:
47.17g - 46.35g = 0.82g
2. Mass of the water lost:
47.17g - 46.63g = 0.54g
3. percentage of water in the hydrate:
.54/.82 = .6585 x 100 = 65.85%
4. Percent Error & possible explanation for error:
pe = [(65.85 - 36.0) / 36.0] x 100 = 82.92% error
I think our percent error was so high because we did not heat up our compound enough to evaporate enough water, thus we were unable to figure out the correct percentage of water in the hydrate.
5. Find the exact formula of the hydrate from the experimental data:
The photo above shows how I calculated the exact formula of the hydrate from the experimental data. Notice the calculations for molar mass on the far right, and the conversions from grams to moles on the left.
The empirical formula obtained from the hydrate was CuSo4 . 17H2O
Our percent error was very high, so I predict that our coefficient on the water was much, much larger than it should have been. The accepted percent of water in the hydrate was 36%, and since ours was almost 66% water I think that we measured far too much water than there actually was in the hydrate.
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