AP Chemistry Pd. 3
8 September 2010
AP Chemistry Laboratory #1
The objective of this lab is to find the empirical formula of silver oxide.
Trial #1 Trial #2 Trial #3
Mass of crucible and 45.5370 21.5892 22.1983
Mass of crucible, lid, 46.0410 22.0730 22.7105
and silver oxide, g
Mass of crucible, lid, 46.0018 22.0399 22.6731
and silver metal, g
Appearance of Dry, hard, white with White colored, Light gray/white,
product silvery glimmer plastered to crucible sandy texture
Silver Oxide Silver Oxygen
Mass (g) .504 .4648 .0392
Percent Composition 100% 92.22% 7.78%
Moles 4.309 1.2
Ratio & Empirical Formula:
Ag:O2 3.5:1 = 7:2
AgxOy Ag + O2 becomes Ag7O4 Ag + O2
1. To find the mass of silver oxide and the mass of the silver metal product, the law of
conservation of mass must be used to calculate the mass of oxygen that combined with silver.
First, take the mass of the crucible, lid, and silver oxide which was previously measured, and
subtract the mass of the empty crucible and lid.
46.0410 g (crucible, lid, and silver oxide)
- 45.5370 g (crucible and lid)
0.46480g (silver oxide)
Then find the mass of the silver metal product by subtracting the mass of the empty
crucible and lid from the mass of the reacted silver product with the crucible and lid.
46.0018g (crucible, lid, and silver product)
- 45.5370 g (crucible and lid)
0.4648 g (silver product)
The law of conservation of mass states that mass that goes into a reaction must be equally
present in the product. Using this concept, the mass of oxygen in the silver oxide can be found
by taking the previously calculated mass of the silver oxide reactant and subtracting the mass of
the silver product.
.504 g (silver oxide)
- .4648g (silver product)
.0392g (oxygen in silver oxide)
2. The percent composition of silver and oxygen in silver oxide is found by separately
taking the mass of the silver and oxygen and dividing it by the mass of the silver oxide. To
convert the products into a percentage, simply multiply by 100.
3. In order to find the moles of the silver and oxygen gas produced by the reaction, the
amount of oxygen and silver used in the reaction must converted into molar mass through
4. Finding the ratio between the number of moles of silver and the number of moles of
oxygen is imperative to finding the empirical formula of silver oxide. To find this ratio, the
moles of oxygen and silver must be divided by the element with the smallest amount of moles in
the reaction, which in this case is oxygen.
Therefore, Ag:O 1.75:1 = 7:2
After acquiring the ratio of silver to oxygen, it is simple to apply this ratio to an empirical
AgxOy Ag + O2 becomes Ag7O4 Ag + O2
5. The above reaction is not balanced. Both sides of the reaction must have the same
amount of silver and oxygen in order for it to be balanced. Because there are 7 silver on the
reactant side, a coefficient of 7 must be added to the product side of the reaction as well. The
oxygen is also imbalanced, as there are 4 oxygen on the reactant side but only 2 on the product
side. The imbalance of oxygen can be corrected by adding a coefficient of 2 to the product side
of the reaction. This balances the equation, as both sides include 7 silver and 4 oxygen.
Ag7O4 7Ag + 2O2
6. The theoretical yield of a product in a chemical reaction reveals the mass of the product
that could be obtained if 100% of the reaction was converted. In this case, the periodic table
must be consulted to find the theoretical empirical formula for silver oxide. Because silver
typically has a charge of 1 and oxygen has a charge of -2, the theoretical empirical formula is
Ag2O. When the molar mass of Ag2O is found, it needs to be converted from the original .504
grams used at the beginning of the experiment into grams silver through the process of
7. The percent yield is used to find the amount of product formed compared to the amount
of product that theoretically could have been obtained. The equation for percent yield is listed
8. Because the percent yield calculated is less than 100%, flaws in the experiment are
exposed. The fact that the percent yield was lower than 100% proves that not enough of the
mass of the product was present in the experiment. This may have been due to the fact that
during the experiment, the silver oxide was being weighed in separate beaker while the crucible
was being heated. When the lead scientist momentarily departed from the laboratory for a
restroom break, the two assistant scientists forgot to reweigh the crucible when the silver oxide
was transferred from the beaker to the crucible, assuming all .504g of silver oxide were
transferred. When the lead scientist returned, some additional calculations were made in an
attempt to compensate for the potential loss of silver oxide (the weight of the empty beaker was
subtracted from the weight of the beaker with some iron oxide remains inside, and then added to
the weight of the crucible, lid, and iron oxide). Despite these additional calculations, the entire
procedure had potential for small errors.
Another possible point of error occurred during the calibration of the scales and the
measuring process. After much weighing had taken place, it became apparent that about five
other scientists were leaning on the table that the scale sat upon and their movements were
affecting some of the measurements. Once this became apparent it was put to a stop, but it was
almost entirely too late at that point in the experiment. The table vibrations could have caused
the silver oxide to be measured at .5 grams but the true mass may have been less. If this was the
case, the amount oxygen in the silver oxide would have been miscalculated because it was
derived directly from the mass of the silver oxide. The mass of the silver oxide would have been
less and thus the silver composing it would have been calculated as less.
Yet another possibility is that not all of the silver oxide got an opportunity to react.
When the final product was examined with a stirring stick, it had a white, shimmery appearance.
The product did not seem extremely silvery until it was broken up into smaller pieces. Because
the silver oxide was not stirred and the flame was hitting only one spot throughout the majority
of the experiment, it is possible that the top layer of the silver oxide was neglected, not getting an
opportunity to fully evaporate and therefore leaving the reaction unfinished. This may have
skewed the ensuing calculations, causing the products of the reaction to be less than the projected
results indicated by the theoretical yield.
Any of these causes or even causes outside of those discussed would have affected the
outcomes of the calculations later performed. These errors would decrease the amount of
product in the reaction, decreasing the numerator in the percent yield equation, leading to a
decreased total percent yield.
The objective of this lab was to initiate a reaction of silver oxide to ultimately calculate
its empirical formula. The possible points of error that were previously discussed may have
caused miscalculations in the determination of even the simplest calculations. If there was less
silver oxide in the crucible before the reaction began due to a measurement problem, all of the
questions would have been solved assuming there was less silver oxide than there actually was.
It would have made an impact as stated above in the question.