Chemistry 113.1 experiment 1. density complete solutions correct answers key

Chemistry 113.1 Experiment 1. Density complete solutions
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Introduction to Chemical Techniques
INTRODUCTION
Density () is defined as the ratio of the mass (m) of a sample to its volume (V):
= m / V
Mass and volume are extensive properties of matter-properties that depend on
the quantity of a
substance. Such properties are not in themselves useful in characterizing or
identifying
substances. Intensive properties such as density however are useful in identifying
substances.
Intensive properties are often determined by taking the ratio of two extensive
properties
measured under constant temperature and pressure conditions. As an intensive
property, density
can be useful in identifying a substance. Density alone cannot absolutely identify
a substance but
can be a useful value contributing to an identification. For example, a colorless
liquid found to
have a density of 1.00 g/mL at 4 0C and 1.0 atmosphere pressure could be water,
since this is the
known density of water. Additional information would be needed to absolutely
identify the
substance. In contrast, a colorless liquid found to have a density of 0.85 g/mL at
4 0C and 1.0
atmosphere pressure could not be water.
In the experiments below, you will use several methods to determine the volume
of samples,
both solid objects and liquids. You will use the electronic balance to determine

the mass of the
samples to 0.001 g (1 milligram). From these measurements, you will determine
the densities of
these samples.
A. DENSITY OF REGULARLY SHAPED OBJECTS
For regularly shaped objects, such as cylinders, the volume can be determined by
measuring the
dimensions of the object with a ruler, then applying the proper formula to
determine the volume.
In this experiment, you will determine the density of a group of objects (all
cylinders)
individually, then by a graphical method.
1. Obtain one set of cylinders from your instructor. Record the CODE on the
container in
your laboratory notebook. Record the color and any other distinguishing
characteristics as
well.
2. Using the plastic ruler in your kit, measure the diameter(d) and the length (or
height-h) of
each cylinder to the nearest 0.5 millimeter (e.g. diameter = 13.5 mm = 1.35 cm).
3. Using the electronic balance assigned to you, determine the mass of each
cylinder to the
nearest milligram (0.001 g).
4. Using the measured diameter and length, calculate the volume of each cylinder
in cm3.
V= r2h = (d/2)2h
Chemistry 113.1
Experiment 1. Density (May 2012)
2
5. Calculate the density r of each cylinder. Report the calculated value for each
cylinder, as
well as the value of the mean (average). [See the addendum regarding the mean
and

mean absolute deviation. These should always be reported whenever three or
more
determinations of the same quantity are the result of identical experiments.]
6. Graph the data for the cylinders with the mass (g) as the y-axis and the volume
(cm3) as
the x-axis. Using the straight line fitting function of the graphing software, find
the
formula corresponding to the best fit of the graphed points to a straight line. The
slope of
this line is the density (Δm/ΔV). Report this value. Compare it to the average
density
reported in (5) above.
B. DENSITY OF IRREGULARLY SHAPED OBJECTS
If an object has and irregular shape, its volume can be determined using
Archimedes’ principle
which states: An insoluble body completely submerged in a fluid displaces its own
volume.
Thus, the volume of the displaced fluid is equal to the volume of the irregularly
shaped object.
1. Obtain a set of mineral samples from your instructor. Record the code
identifying the
sample in your laboratory notebook. Note any distinguishing characteristics of
the
minerals such as color, shape etc.
2. Using the electronic balance record the mass of the samples to the nearest
milligram
(0.001 g).
3. Place approximately 30.0 mL of water in the 100 mL graduated cylinder. Record
the
exact volume to the nearest 0.1 mL.
4. Carefully add the sample or samples to the water in the graduated cylinder
without
causing any water to be lost by splashing. Note: It may be best to determine the

combined volume of 2 or more pieces of mineral together. As long as the
samples are
completely submerged, the greater the increase in volume for the water in the
graduated
cylinder, the more precise the measurement of density will be.
5. Remove the samples from the cylinder, dry them with a paper towel, return
them to the
storage container and return them to your instructor.
6. Calculate the density () of the mineral sample.
C. DENSITY OF LIQUIDS: CONSTRUCTING A CALIBRATION CURVE AND DETERMINING V%
COMPOSITION OF AN UNKNOWN SAMPLE
In this experiment, you will prepare a series of liquid mixtures of known
composition (percent by
volume or V%) and determine their densities. Using your graph of density versus
V% you will
determine the V% of an unknown sample by measuring its density and
comparing it to the
graphed values.
Chemistry 113.1
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1. Using a small beaker, obtain about 10 mL of alcohol (either methanol or
ethanol). Record
the name of the alcohol used.
2. Place a clean dry 10 mL graduated cylinder on the electronic balance and tare
it to 0.000
g.
3. Carefully transfer 2 mL of the alcohol into the graduated cylinder. Read and
record the
exact volume to the nearest 0.1 mL. [ Note: When reading the volume, the level
being
read should be at eye level.]
4. Record the mass of the alcohol. Using the mass and volume, calculate the

density and
enter it in the data table.
5. Add 1 mL of distilled water to the alcohol in the cylinder. Record the exact
total volume
now in the cylinder and total mass. [Note: Total volume should now be about 3.0
mL]
6. Add 2 additional mL of distilled water to the contents of the cylinder (the total
volume at
this point should be approximately 5 mL total). Record the exact total volume
and the
total mass.
7. Add 2 additional mL of distilled water to the contents of the cylinder (the total
volume at
this point should be approximately 7 mL total). Record the exact total volume
and the
total mass.
8. EMPTY the graduated cylinder and dry it.
9. At this point you may need to place the cylinder on the balance and tare to
0.000g once
again. Then-add 2 mL of distilled water to the graduated cylinder and record the
exact
volume and the mass to 0.001 g. Divide the recorded mass of the water by its
volume.
The value should be 1.00 g/mL, which is the known density of water. If this is not
the
case consult with your instructor immediately.
Obtain a sample of alcohol/water of unknown (to you) V% composition .
Determine the volume
and mass of two individual 3 to 4 mL portions (known as aliquots) of this sample.
Record the
exact volume and mass of these two aliquots.
Chemistry 113.1

4
RECORDING DATA
In your laboratory notebook, you should record all data in a format similar to the
suggested
formats below:
A. DENSITY OF REGULARLY SHAPED OBJECTS
Sample Code __________________
Description:
Table A. The dimensions and mass of each object.
Object Diameter (cm) Length or Height (cm) Mass (g)
A
B
C
D
Note: In your laboratory report, this data should be transcribed into a neatly
typed table. A final
column should be added giving the calculated value of density , in g/cm3.
Using the tabulated values above prepare a graph of mass(g) (y-axis) versus
volume (cm3) (xaxis).
A straight line fit has a slope equal to the density of the sample [slope = Δm/ΔV].
Report
this value of the determined density. Using the table of Materials and their
Densities provided,
can you identify the material ?
B. DENSITY OF IRREGULARLY SHAPED OBJECTS
Table B. The mass (or combined masses) of the mineral sample(s), the initial
volume of water in
the graduated cylinder [Vi (H2O)], the final volume after addition of the sample(s)
[Vf (H2O)],
and the volume of the sample(s), Vmineral.
M (g) [Vi (H2O)] (mL) [Vf (H2O)] (mL) Vmineral (mL)

column should be added giving the calculated value of density , in g/cm3. [Recall
that 1 mL=1
cm3]. Using the provided table of mineral densities and descriptions, can you
tentatively identify
the mineral you were assigned ?
Chemistry 113.1
5
C. DENSITY OF LIQUIDS: CONSTRUCTING A CALIBRATION CURVE AND DETERMINING VOLUME
PERCENT (V%) COMPOSITION OF AN UNKNOWN SAMPLE
(1) The alcohol used in this experiment was ____________________
(2) The V% for this alcohol is _______________________
[Note: ethyl alcohol (ethanol) is 95% by volume ethanol and 5% water]
Table C1. Volume of alcohol Valcohol , volume of added water Vwater, total volume
Vtotal , total
mass m and volume percent alcohol V%alcohol
Sample Valcohol Vwater added Vtotal mass total V%alcohol
1 At start (only
entry)
2
3
4
5 (water only)
The volume percent (V%) is equal to: 100 x (Valcohol / Vtotal ). For example, if 2.0 mL
of 95V%
ethanol is initially present, then for a Vtotal of 5.0 mL (after addition of 1 and then
2 mL of
water):
V%ethanol = 100 x(0.95 x 2.0) / (5.0) =38%
column should be added giving the calculated value of density , in g/mL for
each value of

V%alcohol.
In your report, you will graph the values of density (g/mL) (y-axis) versus V% (x-
axis) and
provide a straight line fit to the points on the graph.
Table C2. Volume V and mass m of two aliquots of alcohol/water mixture of
unknown V%.
Volume Mass Density (r)
aliquot #1:
aliquot #2:
Using the graph of density (g/mL) (y-axis) versus V% (x-axis) which you have
prepared,
determine and report the V% of each of the two aliquots.
Chemistry 113.1
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A. TABLE OF MATERIALS AND THEIR DENSITIES
Material Density(g/cm3)
aluminum 2.71
teflon 2.20
polyvinyl chloride 1.37
phenolic 1.32
polyurethane 1.23
acrylic 1.17
nylon 1.15
polypropylene 0.90
B. TABLE OF MINERAL DENSITIES AND DESCRIPTIONS
Density
(g/mL) description Probable ID
5.0 metallic, brassy, crystalline FeS2 pyrite
4.94-5.07 metallic gray hematite, a-Fe2O3
2.75-2.79 light blue with white and yellow-brown aquamarine
4.01 dark brown with red highlights Alamndine garnet
3.12 black Schorl tourmaline
2.93 blue with white lapis lazuli

2.83 dark brown with paler gold and white streaks tiger eye, microcrystalline SiO2
with
iron oxide
2.71 purple and hite amethyst, SiO2 with Fe impurities
2.69 light brown with paler gold and white streaks citrine
2.59 light gray with black Labradorite
2.58 turquose (light) with white amazonite
2.53 light blue with white and brown chrysoprase
2.36 royal blue with white veining sodalite, Na4Al3(SiO4)3Cl
Chemistry 113.1
7
LABORATORY REPORT SUGGESTIONS FOR THIS LABORATORY
In addition to the more general instructions posted on Blackboard, here are some
specific tips
with reference to the Density experiments:
Abstract
One sentence defining density
One sentence stating the results of the regular shaped object density
determination
One sentence stating the results of the irregularly shaped object(s)-minerals
density
determination
One sentence or two giving the results of the density of liquids experiments,
including the
volume percent (V%) of the unknown mixture
Introduction
Explain the difference between intensive and extensive properties and why one
is useful
in characterizing materials
Explain density with the equations, defining all variables or symbols used
Experimental
Summarize the procedures for regularly shaped objects

Summarize the procedures for irregularly shaped objects
Summarize the procedures for density of liquids, including the unknown
sample.
Results/Discussion
Tables of data should be neatly transcribed from the data sheets in your
laboratory
notebook, adding where needed values (such as density) calculated from the
data.
Provide a sample calculation of each type of calculated value (but not every
one).
Show the graphs for density of regular shaped objects (A) and volume percent
as a
function of density (C). State the slope for each and what it signifies.
In the appropriate section, provide the tentative ID for the material of the
regularly
shaped objects (cylinders) and explain your reasoning. Do the same for the
irregularly
shaped object(s) with appropriate reasoning. In each case provide an estimate of
the
uncertainty in the reported density values, with explanation.
Report the value of the V% for the unknown alcohol-water mixture, with
uncertainty.
Chemistry 113.1
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ADDENDUM
REPORTING THE MEAN AND MEAN ABSOLUTE DEVIATION
Calculating the Mean density and Mean Absolute Deviation
1) Determine the Mean: Add all numbers and divide by the count (3)
Example: the density of three cylinders, denoted by letters are found to be:
A : 1.6 g/cm3
B : 2.0 g/cm3
C : 1.8 g/cm3

Mean = (1.6 + 1.8 + 2.0)/3 = 1.8 g/cm3
2) Determine deviation of each result from the Mean ( individual value – Mean)
1.6 - 1.8 = - 0.2
1.8 - 1.8 = 0.0
2.0 - 1.8 = + 0.2
3) Eliminate the + or – sign and take the mean of the absolute deviations
Thus the Mean Absolute Deviation is (0.2+0.0 +0.2)/3 =0.13 g/cm3 round to 0.1
g/cm3
Report the density as 1.8 +/- 0.1 g/cm3 [Alternatively, you can use the Excel
formula
=AVEDEV(1.6, 2.0, 1.8) to obtain the result. ]
Chemistry 113.1
Experiment 2. Hydrate Composition (May 2012)
1
I. INTRODUCTION
The law of definite (or multiple) proportions states that when two or more
elements combine to
form a given compound, they do so in fixed proportions by mass. For example,
sodium chloride
contains 39.3% by mass sodium and 60.7% by mass chlorine. In these
experiments, the law of
definite proportions will be used to determine the empirical formulas of
hydrated ionic salts. An
empirical formula expresses the simplest whole number ratio of atoms or units in
a compound.
(For ionic compounds or hydrates, the unit can be a polyatomic anion or water.)
Hydrates are substances formed when water combines chemically in definite
proportions
with an ionic salt, thereby giving a constant ratio of water molecules to the ions
of the salt.
Hydrates are not mixtures, since the water is coordinatively (covalently) bound to
either the

cation or anion or both in the salt. In CuSO4 • 5 H2O, for example, the bonding
involves four
water molecules coordinatively bound to the Cu2+ ion in a square planar structure
and one
molecule of water bound to the sulfate ion by hydrogen bonds. The anhydrous
(without water)
form of a hydrated salt is produced when all the waters of hydration are lost.
Some examples of
hydrates are listed below:
Formula Common name
2 CaSO4 • H2O plaster of Paris
CaSO4 • 2 H2O gypsum
CuSO4 • 5 H2O blue vitriol
MgSO4 • 7 H2O Epsom salt
Na2CO3 • 10 H2O Washing soda
The • in the formula indicates a kind of chemical bond that usually can be easily
broken. For
example, magnesium sulfate heptahydrate can be converted to anhydrous
magnesium sulfate by
heating:
MgSO4 • 7 H2O (s) → MgSO4 (s) + 7 H2O (g) .
In this chemical reaction equation (or chemical equation), the (s) indicates a solid
and the (g)
indicates a gas. In the appendix, more details about this reaction equation will be
given along
with how these equations are balanced and how they can be used to predict
products of reactions.
In this experiment, you will heat various hydrated salts to determine the number
of water
molecules in the salt.
Chemistry 113.1

2
II. PROCEDURE
Obtain from the instructor a hydrated salt chosen from copper sulfate, calcium
sulfate, and
magnesium sulfate. The difference in the mass of the anhydride and the hydrate
will then be
used to determine the mass of water in the hydrate and, therefore, the empirical
formula of the
hydrate. The procedure, which should be performed on two samples of the same
hydrate, is as
follows:
1. Weigh a clean, dry, labeled crucible. Record the weight in your notebook.
2. Introduce about 1 - 2 grams of the pulverized hydrated salt. Note the
appearance and
color of the solid.
3. Weigh the crucible and contents. Record this weight in your laboratory
notebook.
4. Setup a wire triangle on the iron ring over a Bunsen burner. (Ensuring that the
wire
triangle will hold the crucible in an upright position.)
5. Watch the instructor demonstrate how to setup and properly light a Bunsen
burner and
how to turn-off the burner after use. (Be sure to record this in your notebook for
later
referral.)
6. Heat the crucible and contents in the hottest part of the flame for 5 - 10
minutes. (The
bottom of the crucible should turn a dull red during heating.) Initially, the hydrate
should
be heated slowly by waving the burner flame fairly rapidly under the crucible. If
the
material begins to boil or crackle, the heating is too intense and splattering may
occur.

Within approximately 1 minute, the material should become drier and stronger
heat can
be applied. At the end of the 5 - 10 minute period of heating, allow the crucible
to cool
slightly before transfer.
7. Using clean crucible tongs, transfer the crucible to a desiccator and allow the
crucible to
cool to room temperature.
8. When cool, weigh the dish and the anhydride and record this weight in your
notebook.
9. Heat the crucible in the flame again for 5 minutes, place in desiccator and
allow the
crucible to cool. Once cool, weigh the sample again. Continue the
heat/cool/weigh cycle
until the mass of the sample remains constant. Be sure to record all of your
measurements in your notebook.
10. Place a thermometer in the anhydrous salt and record the temperature.
11. Add a few drops of water to the anhydrous salt near the thermometer and
record the
temperature (once the temperature has stopped increasing). This temperature
change
represents the change from an anhydrous salt to a hydrated salt.
Remember: Write the experimental procedures that YOU followed while you
were doing
the experiment. Be sure to note if the salt splattered or popped out of the
crucible while
heating. These types of observations will be important when discussing
sources of
experimental error.
Chemistry 113.1
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Useful information that should be recorded in the notebook at some point during
the experiment:
Name of salt and formula

Qualitative description of the salt before and after heating.
Temperature of the salt before and after the addition of water
Table 4.1. Masses m in grams (g) necessary to determine the composition of the
salt from the
first trial.
Object m (g) Notes
Clean, dry crucible
Crucible with salt (before
heating)
Crucible with salt after 1st
heat/cool cycle
Crucible with salt after 2nd
heat/cool cycle
Crucible with salt after 3rd
heat/cool cycle (or until stable)
Crucible with salt after 4th
Chemistry 113.1
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Table 4.2. Masses m in grams (g) necessary to determine the composition of the
salt from the
second trial.
Object m (g) Notes
Clean, dry crucible
Crucible with salt (before
heating)
Crucible with salt after 1st
heat/cool cycle
Crucible with salt after 2nd
heat/cool cycle

Crucible with salt after 3rd
Chemistry 113.1
5
III. POST-LABORATORY DISCUSSION AND QUESTIONS
The mass of a single atom is difficult to measure. (For instance, the mass of a
single hydrogen
cation (or proton) is 1.67 × 10-24 g.) Therefore, the mole is defined as the
number of 12C atoms in
exactly 12 grams of 12C. Moreover, the basic unit of mass for elemental chemistry,
namely the
atomic mass unit (amu or dalton) is defined as 1 amu ≡ 1/12 the mass of an
atom of 12C = 1.6605
× 10-24 g. Thus,
The constant 6.022 × 1023 atoms (or molecules)/mole is known as Avogadro's
number NA. Since
the mole and the atomic mass unit are defined using the same scale, 1 amu ×
NA = 1 g/mole.
Thus, the masses given on the periodic table can also be expressed as the
number of grams of the
element per mole of element. The molar mass M of a compound is obtained by
summing the
mass of all of the elements in a compound and, therefore, has units of g/mol.
Moreover, the
definition of a mole when combined with the law of definite proportions implies
that a sample of
H2O will have 2 moles of atomic hydrogen for every 1 mole of atomic oxygen,

while a sample of
MgF2 has a mole ratio of 1 mole of magnesium for 2 moles of atomic fluorine.
Please note that moles are used as the universal conversion factor in chemistry.
The
chemical reaction equation written in the introduction, for instance can now be
read as follows:
1 mole of solid magnesium sulfate heptahydrate decomposes with heating to
generate one
mole of solid magnesium sulfate and 7 moles of gaseous water
Thus, please become more familiar with this difficult concept by reading about
the mole in a
standard freshman chemistry book or on Wikipedia. Also, in the future, when in
doubt,
convert to moles! The post-laboratory questions below will help guide you in the
conversions
between mass in grams and moles of compound and will show you how this
allows you to
determine an empirical formula for a hydrate.
Chemistry 113.1
6
POST-LABORATORY QUESTIONS
Where appropriate, these questions should be answered for each trial.
Remember, do not
write Question 1 and then an answer. Also remember to show all work for the
calculations
for one of the trials.
1. Determine the mass mh of the hydrated salt by subtracting the mass of the
empty crucible
from the mass of the salt and the crucible before heating.
2. Determine the mass ma of the anhydrous salt by subtracting the mass of the
empty
crucible from the mass of the salt and the crucible after the heat/cool cycles are

complete.
3. The difference in the mass of the hydrated salt and the anhydrous salt is the
mass of water
present in the sample. Why is the mass different (i.e., what happened to the
water)?
4. Calculate the molar mass of water.
5. Calculate the moles of water in the hydrated salt by dividing the mass of water
by the
molar mass of water.
6. Calculate the molar mass of the anhydrated salt. (To do this, use the chemical
formula
you wrote from the name of the compound that you used.)
7. Calculate the moles of the anhydride in the sample by dividing the mass of the
anhydride
by the molar mass of the anhydride.
8. Determine how many moles of water are associated with a single mole of
anhydride by
dividing the moles of water by the moles of anhydride. What is the average value
for this
ratio?
9. Using the information, write the formula of the hydrated salt in the form
Anhydride • x H2O , where x is the average value obtained in question 8
10. Is x an integer to the correct precision? If not, why? What sources of error
could have
caused x not to be an integer?
11. Lookup your salt on Wikipedia. Is x in Question 9 an appropriate value based
on the
possible hydrates that your salt can form? What is the percent error in your value,
assuming that the information on Wikipedia for the hydrate is correct?
Chemistry 113.1
Experiment 3. Precipitation reactions (May 2012)
1

I. INTRODUCTION
In Experiment 3, you applied heat from a Bunsen burner to decompose a hydrate
into an
anhydrous salt and gaseous water. A decomposition reaction is one of four
broader categories of
chemical reactions. The remaining categories are precipitation reactions,
acid/base reactions,
and oxidation/reduction reactions. In this experiment, you will investigate
precipitation
reactions. In a precipitation reaction, two aqueous solutions of soluble salts are
mixed and yield
an aqueous solution of a soluble salt and a solid compound. The formation of
the solid is called
precipitation and the solid is called the precipitant.
When ionic compounds dissolve in water, the water interacts with the cation and
anion to
weaken the Coulombic interaction holding the two ions together as a solid. Thus,
as the ions
break apart because of water surrounding the individual ions in the compound,
the solid
dissolves. For example, when CuSO4 dissolves in water, the chemical reaction
equation is
CuSO4 (s) → Cu2+ (aq) + SO4
2- (aq) .
Again, the (s) stands for solid, while the (aq) stands for aqueous. Precipitation
occurs when
aqueous cations and anions form Coulombic interactions that are strong enough
to overcome the
interaction of the water molecules with the separate ions in solution. In other
words, the solid is
more stable than an aqueous solution containing the two aqueous ions.
Compounds that do not
dissolve in water are called insoluble, while those that do are called soluble. Table

5.1 gives
some basic rules for the solubility of ionic salts. These rules should be
memorized, although in
Chemistry 114, you will learn that these rules are not as black and white as they
are presented
here.
Table 5.1. Rules for determining the solubility of ionic compounds
1. Compounds of the alkali metal ions (i.e., the column to the far left on the periodic table) are
soluble.
2. Compounds containing ammonium ion are soluble.
3. Nitrates, chlorates, perchlorates, and acetates are soluble.
4. Chlorides, bromides, iodides are soluble except when combined with lead(II), silver(I) and
mercury(I)
cations. Mercury(II) iodide is also insoluble.
5. All sulfates are soluble except when combined with strontium, barium, calcium, lead(II),
mercury(I), and
silver cations. Small amounts of calcium, silver, and mercury(I) sulfates will dissolve in solution (i.e.,
slightly soluble).
6. Carbonates, phosphates, oxalates, and chromates are insoluble unless they fall under the
categories 1 and 2
(i.e., alkali metal or ammonium ion salts).
7. Sulfides are insoluble unless they fall under categories 1 and 2. Alkaline earth metals (i.e.,
calcium,
strontium, and barium) form slightly soluble sulfides.
8. Hydroxides and oxides are insoluble except for those that fall under categories 1 and 2.
Alkaline earth
metals form slightly soluble hydroxide and oxide salts.
Chemistry 113.1
2
In a precipitation reaction, two solutions containing soluble ionic salts are mixed.
However,
some of the ions, when the solutions are combined, can interact to form
insoluble salts. When
this occurs, a solid forms (i.e., the precipitant) and produces a cloudy solution and
often falls to

the bottom of the container. For example,
Solution A: Aqueous copper sulfate: CuSO4 (aq) or Cu2+ (aq) + SO4
2- (aq)
Solution B: Aqueous sodium carbonate: Na2CO3 (aq) or 2 Na+ (aq) + CO3
2- (aq)
Notice that we can represent aqueous solutions (aq) in two ways, namely one
with the ionic
compound formula (aq) and one with the individual ions (aq). Both ways are
equally valid.
Also notice that when sodium carbonate goes into solution, we obtain one mole
of carbonate
anion for every two moles of sodium cations. Bulk solutions can have no net
charge. Thus,
when you sum the charges for all ions in solution, you must obtain zero. The
integer number to
the left of the sodium cation insures that the solution remains at zero net (or
total) charge. Then,
when the two solutions are mixed, we have the following chemical reaction
equation:
Cu2+ (aq) + SO4
2- (aq) + 2 Na+ (aq) + CO3
2- (aq) → CuCO3 (s) + 2 Na+ (aq) + SO4
2- (aq)
The equation written above is called an ionic reaction equation. Notice that the
sodium cation
and the sulfate anion appear on both sides of the equation. This is because he
combination
Na2SO4 is soluble. These ions are called spectator ions because they do not
participate in the
chemical reaction that yields the solid copper(II) carbonate. Thus, the net ionic
reaction is
Cu2+ (aq) + CO3

2- (aq) → CuCO3 (s)
Notice how both of these reaction equations have the same number of each
atom on the right
hand and left hand side of the equation and that both of these equations also
have the same net
charge on both sides of the equation. Insuring that the charge and mass on both
sides of a
chemical reaction equation is the same are both part of balancing a chemical
reaction equation.
II. EXPERIMENT
1. Given the following aqueous solutions, predict which mixtures of any two of
these
solutions would yield precipitants and determine the formula for the precipitant:
a. Copper(II) sulfate b. Barium nitrate
c. Sodium chloride d. Silver nitrate
e. Lead nitrate f. Sodium sulfate
2. Show your predictions to the instructor.
3. Test to see if your predictions where correct by placing a small amount of the
appropriate
solution into a clean test tube and adding the second solution. You must test all
of the
reactions that you predict would form a precipitant. When the precipitant forms,
record
the color of the precipitant. You must also test at least two reactions that, based
on Table
5.1, you predict would not form a precipitant.
Chemistry 113.1
3
III. Post-laboratory discussion and questions
You must always balance chemical reactions before using these reactions to
predict the amount

of product one might expect from the reaction (i.e., the theoretical yield of the
reaction). In this
experiment, one has probed solubility and, therefore, must now learn how to
write a chemical
reaction equation and balance it. When balancing a chemical reaction equation,
the mole ratio
of various compounds in the reaction can be adjusted to balance both mass and
charge for the
reaction. However, the formula of a compound or a polyatomic ion does not
change. Therefore,
do not change atomic symbols or subscripts in chemical formulas while
attempting to balance
reactions. Some basic guidelines are
1. Balance atoms other than H and O first
2. Pure Elements [e.g. Fe (s) or Cl2 (g) ] should be balanced last
3. Balance as a unit any polyatomic ions that appear unchanged on both sides of
the arrow.
As an example, let us balance the reaction of aqueous copper sulfate with
aqueous sodium
hydroxide. The first step is to write the ionic reaction. Thus,
CuSO4 (aq) + NaOH (aq) → products
To determine the products, we need to look at the solubility rules in Table 5.1.
Rule 8 states that
hydroxides are insoluble unless that are formed with alkali metal cations or with
ammonium.
Since copper is not an alkali metal, this tells us that copper hydroxide will be
insoluble in water.
Thus, one of the products will be copper(II) hydroxide. We also know from Rule 5
that almost
all sulfate salts are soluble. Thus, sodium sulfate will not form an insoluble salt.
Therefore, the
sodium ion and the sulfate ion are spectator ions in this reaction and can be

removed from both
sides of the equation. Using this information, we can now write the unbalanced
net ionic
reaction
Cu2+ (aq) + OH- (aq) → Cu(OH)2 (s)
To balance the reaction, we can see that we have one mole of Cu2+ on the left
side of the
equation and one mole of Cu2+ on the right side of the equation. However, we
only have one
mole of OH- on the left side of the equation and two moles on the right side of
the equation.
Therefore, we need to place a 2 in front of the OH- on the left side of the
equation. Once we
have done so, we obtain the balanced reaction (check the net charge on both
sides)
Cu2+ (aq) + 2 OH- (aq) → Cu(OH)2 (s)
What happens if one mixes aqueous sodium sulfate with aqueous potassium
nitrate. Since
potassium sulfate is soluble (Rule 5) and sodium nitrate is soluble (Rule 3), no
reaction occurs!
Thus, one would write K2SO4 (aq) + NaNO3 (aq) → no reaction.
Chemistry 113.1
4
POST-LABORATORY QUESTIONS
1. Write the balanced chemical reaction equations for all reactions that yielded a
precipitate.
2. Did the color of the solution and the color of the precipitate match? Speculate
why or
why not. (This is a very complicated question. We are not expecting four pages to
answer this question after extensive research. We are asking that you think about
what

might cause a color difference.)
3. Did any of the solutions that you expected should not have yielded a
precipitate actually
yield a precipitate? If so, state why you think this may have happened ?

Chemistry 113.1 experiment 1. density complete solutions correct answers key

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Chemistry 113.1 experiment 1. density complete solutions correct answers key