Lab 3 A & B - Measuring the Composition of Matter – Percent Water in a Hydrate &
Identification of an Unknown Hydrate
Summer 2015
Chemistry B1A / by Mike Daniel and Kenward Vaughan
The Law of Definite Proportions (also called Law of Constant Composition) is a summary of observations that a
compound will always have the same proportions by mass. For example, water will always be composed of 11.19%
hydrogen and 88.81% oxygen by mass. This is true regardless of where the water came from. A mixture can have
variable proportions while a substance always has the same ratio of elements by mass.
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Another important law is the Law of Conservation of Mass . When a change occurs in matter, the mass of all the
materials involved at the start are still present after the change occurs. When water melts, all the mass is still there
2
after it has liquefied. When 16.0 grams of methane react completely with 64.0 grams of oxygen, the products’
masses total 80.0 grams. There is no mass lost or gained. Analysis by mass is called gravimetric analysis.
Any theory of matter must conform with these (and other observed) laws.
This lab has you explore a the chemical change when you heat a hydrate, the results of which depend on the above
laws. Try to determine where/how they fit into your experimental results.
First Lab Period - Procedure Part 3A
1. Qualitative tests of a CuSO4 hydrated salt. Record all data and observations in your lab notebook. You need
to add enough solid CuSO4 hydrate to a large test tube to get about ½” (1.27 cm) deep.
A. In this part of the experiment, you are determining the decomposition products of a hydrate.
B. Record the purpose of this lab and your initial observations about the copper (II) sulfate hydrated salt.
C. You could record the following steps observations in your lab notebook in a table similar to the following.
Table 1 – Model Observation Table for copper (II) sulfate hydrate test
Hexane
2-propanol
Liquid
Observations
Observations with
cobalt paper
Observations
from mixing
compounds with
CuSO4 hydrate
1.
2.
1
2
Methylene
chloride
toluene
Water
Liquid from
step D2
NA
Place about one mL (about ½” or a small finger width) each of hexane, 2-propanol, methylene
chloride, toluene, and water, into separate small, dry test tubes. Record your observations. Note
clear is not a color. Be sure the tubes are labeled in some way. You will need 6 pieces of blue
cobalt paper. Do not touch the paper with your fingers. Use a watch glass and forceps to get the
paper.
Test each liquid by putting one drop, (using a stirring rod) on a piece of cobalt paper. Record your
observations.
More properly, this should be the Law of Conservation of Energy and Mass, but we will discuss this later.
“React completely” implies there are no leftover reactants.
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3.
Use your spatula to place a small amount (about the size of a small match head) of the copper (II)
sulfate hydrate into each of these liquids. It is better to have too small an amount than to have
too much of the solid. Use a vortex mixer to stir them.
4. Record your observations. Observation should include colors (or colorless if it has no color) of the
mixture, is it homogeneous or heterogeneous, and what state(s) of matter are present such as
solid, liquid or gas.
5. Dispose of the mixtures in the appropriate waste container in the hood room.
D. Slowly add heat to the solid copper (II) sulfate remaining in the large test tube over a Bunsen burner using
a test tube holder. Heat only the bottom part of the tube. This is best done by holding the test tube
almost sideways and passing the bottom back and forth through the flame of a Bunsen burner.
1. Record your observations.
2. If you see a liquid form in the test tube, test the liquid with blue cobalt paper. Record all of your
observations in your lab notebook table.
3. Heat the test tube until the solid is mostly a faint blue color. Stop heating if the solid starts to
turn brown or black. PLEASE NOTE WE ARE HEATING, NOT BURNING THE SAMPLE.
4. Allow the tube to cool completely.
5. Record your observations.
E. While holding the bottom of the tube with your hand, slowly add about 1 mL of water into the salt in the
bottom of the tube. Record your observations.
2. Quantitative (Gravimetric) tests of Epsom Salt
We will now quantitatively measure the amount of decomposition products resulting from heating a Magnesium
Sulfate hydrate called Epsom Salt. In this part of the experiment you are determining the mass percent of the
gaseous decomposition product from heating the Epsom Salt. Using your observations from above, you can
calculate the number of moles of the two decomposition products. The next page shows a model example of a
data table with calculations. You need this in your lab notebook substituting MgSO 4 for the copper (II) sulfate. Be
sure you understand the calculations and significant figures in this model table.
Weigh about 1.5 g of this hydrated compound to the 0.001 g. When it says about 1.5 g, any amount from 1.4 – 1.6
would work, recorded to the 0.001 g. You will need to heat this compound to decompose it until it has reached a
constant mass. This is not evaporation of water or burning the compound, but decomposing the hydrate. Heat it in
a large test tube vigorously for 15 minutes, cool, weigh, and repeat until the change in mass of the anhydrous salt
from one heat cycle to the next is less than 0.01 g. When finished heating the Epsom salt, calculate the percent by
mass of the compound that was water and the percent by mass that was MgSO 4. Note the mass percent of water
theoretically should be a little greater than the mass percent of anhydrous salt. Check your calculations to be sure
that the per cent water is a little greater than the mass percent of the anhydrous salt.
Using the molar masses of the elements and the experimental masses of your compound before and after heating
it, you can determine the experimental mole ratio of water to magnesium sulfate. These calculations should be
clearly labeled and neatly written in your lab notebook. The average mass of one mole of hydrogen atoms (to 0.01)
is 1.01 g / mol H and the average mass of one mole of oxygen atoms is 16.00 g / mole O. Adding two moles of
hydrogen plus one mole of oxygen will give a mass of 18.02 g/mole H2O. This is called the molar mass, or the mass
of one mole of water. Calculate the following (in your lab notebook) to determine the formula of magnesium
sulfate hydrate, (i.e., calculate the integer n in the formula MgSO4 n H2O).
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Table 2 – Calculations of water to anhydrous magnesium sulfate mole ratio.
𝑚𝑜𝑙 𝐻2𝑂
Moles of H2O lost from heating Epsom Salt=
Exp g H2O X
Moles of MgSO4 = (the mass remaining after heating the
Epsom Salt)
n=
n rounded to the nearest integer
Exp g MgSO4 X
18.02 𝑔 𝐻2𝑂
𝑚𝑜𝑙 𝑀𝑔𝑆𝑂4
𝑥 𝑔 𝑀𝑔𝑆𝑂4
moles H2O / moles MgSO4
n=
The value of n should be within 0.1 of an integer. The CRC Handbook of Chemistry and Physics, located in the lab,
has the formula for Epsom salt (magnesium sulfate heptahydrate) listed alphabetically under MgSO4 in the section,
Table 3 – Model decomposition data table and calculations
1. Mass of empty test tube
2. Mass of test tube + copper (II) sulfate hydrate – CuSO4 n H2O
3. Mass of CuSO4 n H2O - (#2 - #1)
st
4. Mass after 1 heating= TT mass + heated hydrate
5. Mass of after 2nd heating step=TT mass + heated hydrate
st
nd
6. Percent of #3 lost going from 1 to 2 heating
(#4-#5)/#3 X 100% - Show work
88.577 − 88.575
|
| 𝑥100%
1.502
Mass / g
87.615
89.117
1.502
88.577
88.575
0.1%
7. Total Mass Lost = #2 - #5 = (TT + hydrate) – (TT+anhydrous cmp) =
Mass of water lost
0.542
8. Mass remaining = Mass of anhydrous copper (II) sulfate - CuSO4 =
#5- #1
0.960
9. Mass Percent of water in CuSO4 n H2O – Show Work
#7
𝑥 100% =
#3
36.1%
10. Moles of Water lost - Show Work
1 𝑚𝑜𝑙 𝐻2𝑂
=#7 x
=
0.0301 mol
11. Moles of anhydrous copper (II) sulfate – Show Work
1 𝑚𝑜𝑙 𝐶𝑢𝑆𝑂4
=#8 x
=
0.00602 mol
𝑥 𝑔 𝐻2𝑂
𝑥 𝑔 𝐶𝑢𝑆𝑂4
12. Ratio of moles H2O to moles CuSO4 – Show Work
#10
=
=
#11
13. Formula of copper (II) sulfate hydrate = CuSO4 n H2O – Use #12 to
determine an integer for n.
Lab_3_AB_HydrateConstantComp_S2015
5.00 𝑚𝑜𝑙𝑒𝑠 𝐻2𝑂
1 𝑚𝑜𝑙𝑒𝑠 𝐶𝑢𝑆𝑂4
CuSO4 5H2O
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Physical Constants of Inorganic Compounds and has the molar mass of Epsom salt. You need to find this to
complete the analysis in the report. The theoretical mass percent of water in MgSO4 n H2O will be
𝑛 𝑥 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝐻2𝑂
𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝑀𝑔𝑆𝑂4 𝑛𝐻2𝑂
𝑥 100% = Mass % of Water in Epsom Salt
You will compare this theoretical mass percent to your experiment value by calculating an experimental error. This
is a measure of how accurate your data is.
Experimental error = |
𝑡ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙−𝑒𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙
𝑡ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙
| 𝑥 100%
When doing the above calculation, remember that addition and subtraction have different rules for determining
the number of significant figures than when doing multiplication and division. Do not round your numbers to the
proper number of significant figures till you finish the calculation. The 100% has an infinite number of significant
figures. You should finish your calculations in your lab notebook and review them with the instructor before
leaving lab.
Second Lab Period - Lab 3B Parts 3 & 4 - Quantitative (Gravimetric) and Qualitative Analysis of an unknown
hydrate
3. Gravimetric Analysis - Record all of your observations and data in your lab notebook.
There are 12 containers of unknown hydrates in the lab. You will be assigned an unknown hydrate. Record its label.
Weigh as precisely as possible (i.e., to the nearest 0.001 g) about 1.5 g. You will heat it similar to how you heated
the Epsom salt to remove the waters of hydration. By comparing the weight percent of water from all the
unknowns from other students’ data, determine which of the 12 hydrates are the same compound. You should be
able to group each of the twelve unknowns with other unknowns. Check with your instructor to validate your mass
data before listing it on the board. Your answer must be within acceptable limits. If your answer is not within
acceptable limits you will need to check your calculations, and if you think you have done your calculations
correctly, you will need to repeat the work.
4. Qualitative analysis of an unknown hydrate
Record all observations in your lab notebook. Your observations will help you later in the semester identify an
unknown compound. Use the following test to identify your unknown from above. The flame emission test and the
pH test will be demonstrated to students. The following table list properties of some compounds.
Magnesium sulfate, MgSO4
Magnesium does not emit a visible flame, but sodium impurity may
emit a small orange flame. The Mg cation and sulfate anion have a
neutral pH (~7) in water.
Sodium carbonate, Na2CO3
Sodium heated in flame emits an intense orange flame. The
carbonate makes this compound more basic than magnesium
sulfate (i.e., pH>7) but less basic than phosphates. Carbonates and
bicarbonates (HCO3 ) react with acid to produce CO2 bubbles.
Sodium phosphate, Na3PO4
Sodium heated in flame emits an orange flame. The phosphate
makes this compound more basic than sodium carbonate, (i.e., pH >
pH of sodium carbonate and greater than 7)
Potassium phosphate, K3PO4
Potassium heated in flame emits a violet flame. (May have some
orange flame due to sodium impurity). Basic pH (pH>pH of sodium
carbonate and > 7).
Record your observations from the following test. Get a small amount (like the head of a large match) of your
unknown on a watch glass.
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1. Flame emission test - Create a loop at the end of a Nichrome wire and clean the loop by heating it in the tip of
the inner Bunsen burner flame till it no longer emits a orange-yellow light. Stick the hot wire into your unknown
powder and then put this solid on wire in the Bunsen burner flame. Observe the color. With the potassium
compound you might first see the orange-yellow due to sodium contamination before your see the potassium
flame color.
2. pH test - Add a drop of water to your unknown and test with wide range pH paper. Compare the color of the
paper to the color scale on the dispenser. For comparison you could also measure the pH of distilled water, an acid
(1 M HCl) and a base (1 M NaOH). These last 3 measurements are without mixing the unknown with them.
3. Carbonate/Bicarbonate test - Add 1 M HCl (hydrochloric acid) to your unknown and record your observations.
Your observations from the above test should help you identify the unknown salt in this lab and in a later lab. This
identification will be included in your lab 3 report. Use the mass of water lost and the anhydrate salt mass from
your section 1.3 results to determine the number of waters of hydration relative to one of the anhydrous salt’s
.
formula unit. (for example, in CoCl2 6H2O, the ratio is 6 / 1). Look up your compound in the Handbook of Chemistry
and Physics, to see the accepted value of your hydrate and calculate a percent error. For example if you calculate a
ratio of 6.2 and the accepted value to 6 you would calculate.
% Error = |
exp 𝑟𝑎𝑡𝑖𝑜−𝑎𝑐𝑐𝑒𝑝𝑡𝑒𝑑 𝑟𝑎𝑡𝑖𝑜
𝑎𝑐𝑐𝑒𝑝𝑡𝑒𝑑 𝑟𝑎𝑡𝑖𝑜
| 𝑋 100% = |
6.2 − 6
6
| 𝑋 100% = 3%
Note above the 6 and 100% both have infinite significant figures and the 6.2 has 2, resulting in the numerator (the
top of the fraction), having 1 significant figure, and the quotient having 1 significant figure.
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Lab 3 – Report
Full Name____________________________
Due Date_________________________
Partner’s Name__________________________
Answer the following questions using complete sentences. Attach your lab notebook pages following these pages.
1. In all the quantitative work, what is the point of heating to a constant mass?
2. What do you conclude about the identity of the gas/liquid produced by heating the copper (II) sulfate hydrate?
Use all of your observations to support your conclusions.
3. In a chemical equation, energy added to a reaction can be shown as a reactant, and energy given off by a
reaction can be shown as a product. Using E as a symbol for energy, write a chemical equation for the reaction of
copper (II) sulfate hydrate in step 1.D. Write a second equation for step 1.E.
4. Fill in the page numbers of your Epsom salt (Part 2 )data table and calculations. ___________ and __________
Fill in the blank with your calculated result, and your attached lab notes should be clearly labeled and clearly show
the following calculations.
 _____________________the experimental mass percent of water in MgSO4 n H2O.







_____________________the experimental number of moles water and moles of MgSO 4
_____________________Your experimentally determined value of n rounded to the correct significant
figure for MgSO4 n H2O.
___________________Your calculated value for the theoretical mass percent of water in MgSO4 n H2O.
_________________Your experimental error for the mass percent of water in MgSO4 n H2O.
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5. Following your lab notes, attach a table of mass percent of water in the unknown hydrates (part 3). Below,
group the hydrates into groups of the same compound. What is the basis for deciding whether or not two
unknowns are the same compound?
6. Summarize your observations and identify your unknown compound. Calculate an experimental error for the
amount of water in your hydrate compared to the accepted value.
Unknown Label and formula__________________ = _________________________
7. Specifically where are precision and accuracy used in this lab?
8. Discuss how the Laws of Constant Composition and Conservation of Mass are used in this experiment
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