Download Laboratory 3

Laboratory
3
Development of an Equation
Objectives
•
Apply laboratory procedures and make observations to investigate a
chemical reaction. Based on these observations, identify the pattern of
reactivity and communicate what occurred in the reaction by writing a
chemical equation and a particle diagram consistent with the data.
Introduction
Suppose two substances are added together and they begin burning, with smoke
ULVLQJDQGÀDPHVMXPSLQJ7RDFKHPLVWWKLVLVMXVWDQRWKHUGD\LQWKHODERUDWRU\
(and a great demonstration!), and she may describe this phenomenon in different
ways. One way is to record macroscopic observations. What took place? What
was measured? What was seen, heard, smelled? Or perhaps the chemist thinks
about her observations and infers (concludes based on evidence and reasoning)
what happened to the particles. What did the atoms, ions, or molecules do as they
LQWHUDFWHGZLWKHDFKRWKHUDQGKRZGLGWKHVHLQWHUDFWLRQVSURGXFHD¿UH"$FKHPist cannot actually see the atoms, but based on the experimental evidence it might
be possible to provide an explanation, a theory, that is consistent with the observations. Finally, the chemist may decide to write a chemical equation to symbolize
what took place during the reaction. All of these are legitimate ways to describe
what transpired and in this experiment, you will utilize each one.
Chemical equations are usually presented as “facts” in textbooks, but how does a
chemist know that the reaction really took place as written? Where is the evidence?
What led to the inference? In this experiment, you will investigate a particular
13
Laboratory 3
reaction, make observations, infer what you think happened to the atoms and ions, and
¿QDOO\SURSRVHDFKHPLFDOHTXDWLRQFRQVLVWHQWZLWK\RXUGDWD
Discussion
Chemical Equations and Particle Diagrams
Chemical equations are used to describe a chemical reaction with symbols. For example,
nitrogen, N2, and oxygen, O2, can combine to form nitric oxide, NO. The chemical equation for this is:
N2(g) + O2(g) $ 2 NO(g)
Reactants
Product
In this equation, the (+) symbol indicates that nitrogen reacts with oxygen and the arrow
indicates that nitric oxide is formed. The chemical formulas on the left side of the equation
are collectively known as the reactants and those on the right side as the products. In
this case we have one kind of product, NO(g). The letter “g” in parentheses is included to
indicate that these are in the gaseous phase. Other phase symbols include (s) for solid, (l)
for liquid, and (aq) for aqueous or in water. The subscripts in the formula tell us how many
of each atom there are in a molecule. In the formula above, there are two nitrogen atoms
in the gas molecule N2, two oxygen atoms in the gas molecule O2, and one nitrogen atom
with one oxygen atom in the gas molecule NO.
The number in front of the chemical formula, or FRHI¿FLHQWV, indicate how many molecules
are present. In a chemical reaction atoms are not created or destroyed, thus there should
be an equal number of atoms of each element on the left and right sides of the equation.
An equation is balanced when there are an equal number of atoms of each element on the
left and right sides of the equation. In the equation above, we can see that one nitrogen gas
molecule and one oxygen gas molecule combine to form two nitric oxide gas molecules.
There are a total of two nitrogen atoms on the left side of the equation and two nitrogen
atoms on the right side of the equation. The same is true for the oxygen atoms, so this
chemical equation is balanced. This same reaction may also be described by indicating
what happens to the particles:
14
Before
After
N2(g) + O2(g)
2NO(g)
Development of an Equation
Question 3.1: Nitrogen and oxygen can combine to form other molecules containing N and O. One is nitrous oxide, or N2O. N2O is commonly used in dentistry for
its anesthetic and analgesic effects and is better known as laughing gas.
Write a balanced chemical equation for N2 reacting with O2 to form laughing gas.
In the boxes below, draw a particle diagram showing the reactants (before) and the
products (after) the reaction. The drawing should be consistent with your balanced
equation.
Reactants (Before)
Products (After)
Common Patterns of Chemical Reactivity
Different types of common chemical reactions are discussed in your textbook. These
include:
1. Combination reactions, which describe the reaction of two or more reactants to form
one product.
2. Decomposition reactions, in which one substance reacts to form two or more compounds or elements. This type of reaction typically occurs when compounds are heated.
3. Combustion reactions, which often occur quickly and evolve heat in the form of a
ÀDPH0DQ\RIWKHVHUHDFWLRQVLQYROYHK\GURFDUERQVDQGR[\JHQUHDFWLQJWRIRUPFDUbon dioxide and water.
4. ([FKDQJHPHWDWKHVLVUHDFWLRQV in which the positive and negative ions appear to
exchange partners.
Question 3.2: Is the chemical reaction N2(g) + O2(g) $ 2 NO(g) an example of a
combination, decomposition, combustion, or exchange reaction?
Overview of Experiment
In this experiment you will investigate the reaction that takes place when an aqueous lead
nitrate solution is combined with an aqueous potassium iodide solution. Your objective
is to determine the appropriate pattern of reactivity and propose a chemical equation and
particle diagram that is supported by your laboratory observations.
To get started you must make observations. In Part A you will combine a lead nitrate
solution with a potassium iodide solution and record your observations. What atoms, ions,
15
Laboratory 3
or molecules do you think are in each solution before they are combined? After they are
combined, is there evidence for a chemical reaction? Is there any indication that a gas is
formed?
What happens to the solutions of lead nitrate and potassium iodide when they are combined? Is there a chemical reaction? To answer these questions you must be able to identify
the presence of different ions. In Part B you will perform qualitative tests for the different
possible ions in this experiment. Procedures are given for these qualitative tests. Make
careful observations when the known ions are tested, since you will be repeating these tests
with your actual sample and comparing the results.
In Part C you will apply your understanding of qualitative tests to determine which ions
are in the solid product that forms when lead nitrate and potassium iodide solutions are
mixed.
Determining the identity of ions comprising your solid product is important. After Part C
you probably have enough evidence to select the appropriate pattern of reactivity. However, you need more information to write a chemical equation. The qualitative analysis
determined what ions are present, but it did not provide any quantitative information. A
chemical equation includes quantitative information and communicates the amounts of the
different substances. To get this information you must complete a quantitative test.
In Part D you will perform a quantitative test to determine the exact ratio of ions in the precipitate that formed during your reaction. This will be accomplished by mixing solutions
of lead nitrate and potassium iodide (the reactants) in varying ratios and making observations. There is an exact stoichiometric ratio in which these two substances will combine
with each other such that neither one remains. However, for every other ratio the reaction
will take place, but one of the reactants will be “leftover.” The reactant that remains after
the reaction was “in excess.” The reactant that was completely consumed is the limiting
reactant.
To illustrate the idea of reaction stoichiometry, consider the burning of hydrogen gas in
oxygen to form water vapor:
2 H2(g) + O2(g) $ 2 H2O(g)
In the balanced equation, the H2:O2 ratio is 2:1. When hydrogen gas is combined with
oxygen in a 2:1 ratio, all of the H2 will react with all of the O2 to produce H2O(g) and there
are no “leftover” reactants.
16
Development of an Equation
What if different ratios of H2 to O2 are used? There will be leftover reactants:
Ratio of H2:O2
Before (Reactants)
After (Products)
4H2(g) + O2(g)
2H2O(g) + 2H2(g)
3H2(g) + O2(g)
2H2O(g) + H2(g)
2H2(g) + O2(g)
2H2O(g)
2H2(g) + 2O2(g)
2H2O(g) + O2(g)
2H2(g) + 4O2(g)
2H2O(g) + 3O2(g)
4:1
3:1
2:1
1:1
1:2
Question 3.3: In the table above, circle the leftover reactants. What is the limiting
reactant for each ratio H2:O2?
H2:O2 Ratio
Limiting Reactant
4:1
3:1
2:1
1:1
1:2
17
Laboratory 3
How could a chemist determine the ratio of H2:O2 that does not have leftover reactants?
They need a test to determine whether H2(g) or O2(g) remains. This is actually a simple experiment; if a wooden splint is lit, extinguished, and then placed in a container of hydrogen
gas while still glowing, it produces a loud “pop.” However, if the glowing splint is placed
in a container of oxygen gas, it reignites.
Question 3.4: Suppose H2(g) and O2(g) are combined in the ratios shown in the
previous table to produce H2O. For which ratios of H2:O2 would the splint test
result in a “pop”? For which ratios of H2:O2 would the splint reignite? How does
this analysis lead the chemist to the correct chemical equation?
In Part D you will apply quantitative reasoning to determine the ratio of potassium iodide
reacting with lead nitrate in your chemical equation. This will be accomplished by mixing
drops of each solution in different ratios and then testing the supernatant (the liquid above
the precipitate) for excess ions. Instead of doing a test with a glowing splint, however, you
will determine whether there are excess ions by checking whether a precipitate can be
formed with the ions (that may) be leftover. The solution in which neither ion is in excess
indicates that the ions were added in the correct ratio to exactly form the product. Note that
having neither ion in excess requires careful measurement of amounts of the two solutions.
A very small amount of precipitate will form while testing the supernatant liquid if there
is a slight excess of either reactant. It is therefore important not only to measure carefully,
but also to note the amount of precipitate when testing for the ion in excess. Can you apply
quantitative reasoning from Part D to determine the chemical equation? Try question 3.5.
18
Development of an Equation
Question 3.5: When solutions of barium chloride and sodium sulfate are mixed, a
white precipitate of barium sulfate forms:
Barium chloride + sodium sulfate $ barium sulfate (a white precipitate) + sodium
chloride
7RGHWHUPLQHWKHFKHPLFDOHTXDWLRQWKHIROORZLQJ¿YHWHVWWXEHVZHUHSUHSDUHGDQG
tested:
Drops of Drops
of
Barium Sodium
Chloride Sulfate
Drops
of
Water
Tube 1
6
2
22
Tube 2
6
3
21
Tube 3
6
6
18
Tube 4
6
12
12
Tube 5
6
18
6
Observation
Inference for the
Barium Chloride to
Sodium Sulfate Ratio
White precipitate
formed
White precipitate
formed
White precipitate
formed
White precipitate
formed
White precipitate
formed
A 3:1 ratio produces
barium sulfate.
A 2:1 ratio produces
barium sulfate.
A 1:1 ratio produces
barium sulfate.
A 1:2 ratio produces
barium sulfate.
A 1:3 ratio produces
barium sulfate.
These results are inconclusive and do not narrow down the barium chloride to sodium
sulfate ratio in the balanced chemical equation because every reaction produces barium
sulfate. However, by testing each supernatant (the liquid above each precipitate) it is possible to determine which ions are “leftover” and in excess. In the table below, identify the
ions present in the supernatant (Ba2+, Clí, Na+DQGRU624í) prior to additional barium
chloride or sodium sulfate.
Tube 1
Tube 2
Tube 3
Tube 4
Tube 5
Barium Chloride
Added to Supernatant
Observation
Sodium Sulfate
Added to
Supernatant
Observation
No additional
precipitate formed
No additional
precipitate formed
No additional
precipitate formed
White precipitate
formed
White precipitate
formed
White precipitate
formed
White precipitate
formed
No additional
precipitate formed
No additional
precipitate formed
No additional
precipitate formed
Ions Present in the
Supernatant Prior to
Adding Additional
Reagents
Question 3.6: Use the information in the table above to determine the empirical
formula of barium sulfate. Support your reasoning.
19
Laboratory 3
Question 3.7: Write a balanced chemical equation for the reaction that occurs
when barium chloride and sodium sulfate are mixed.
Materials Required
Equipment
15–4-mL test tubes
medicine dropper
micropipet
micro spatula
water bath (100-mL beaker)
WHVWWXEHUDFN
Bunsen burner, wire gauze
thin-walled rubber tubing
ring stand, ring
Common Equipment
centrifuge
Chemicals
0.1 M lead nitrate, Pb(NO3)2P/VWXGHQW
0.1 M potassium nitrate, KNO3P/VWXGHQW
0SRWDVVLXPLRGLGH.,P/VWXGHQW
3 M nitric acid, HNO3P/VWXGHQW
5% thioacetamide, CH3CSNH2P/VWXGHQW
3% hydrogen peroxide, H2O2P/VWXGHQW
1,2-dichloroethane, C2H4Cl2P/VWXGHQW
Cautions
%HFDUHIXOWRDYRLGEXUQVIURPULQJDQGRSHQÀDPH6LQFHOHDGVDOWVDUHWR[LFVXEVWDQFHV
be sure to wash your hands thoroughly. 1,2-dichloroethane (dichloroethane) is a suspect
carcinogen and may be irritating to the respiratory tract. Heating thioacetamide solutions
produces H2S, an inhalation poison. Do this procedure well within the hood. Fume hoods
must be on.
Procedure
/DEHODFOHDQGU\P/WHVWWXEH³OHDGQLWUDWH´DQG¿OOLWZLWKWKHDSSURSULDWHVROXWLRQ
/DEHODVHFRQGFOHDQGU\P/WHVWWXEH³SRWDVVLXPLRGLGH´DQG¿OOLWZLWKWKHDSSURSULDWH
solution. Always take drops from these tubes.5H¿OODVQHFHVVDU\
A. Initial observations.
In this section of the experiment you will observe the reaction that occurs when solutions
of potassium iodide and lead nitrate are mixed.
1. Add 5 drops of each solution to a 4-mL test tube and observe what happens. It may be
helpful to centrifuge the test tube after the reaction. Record your observations in your
notebook. Save this tube for use in Part C.
B. Qualitative tests for ions.
In Part C we will use chemical tests to determine which ions are present in the product of
WKHDERYHUHDFWLRQ+RZHYHUEHIRUHWHVWLQJRXUSURGXFWZHPXVW¿UVWVHHKRZHDFKRIWKH
anions and cations present in this experiment react. That is the task for this section.
20
Development of an Equation
Procedure for Testing Anions
In this section of the experiment, you will see what the reactions of the anions (iodide and
nitrate) look like when they react with hydrogen peroxide. Be sure to make careful observations in your notebook. Later, you will use this test to determine which anion is present in
a solution by repeating the experiment and comparing your observations.
2. Add 5 drops of KI, 5 drops nitric acid, 10 drops dichloroethane, and 5 drops hydrogen
peroxide to a clean 4-mL test tube. Stopper and shake the test tube. You may decide to
centrifuge the solution in the test tube to produce two distinct layers.
3. Add 5 drops of KNO3, 5 drops nitric acid, 10 drops dichloroethane, and 5 drops hydrogen peroxide to a clean 4-mL test tube. Stopper and shake the test tube. You may decide
to centrifuge the solution in the test tube to produce two distinct layers.
Procedure for Testing Cations
In this section of the experiment you will see what the reactions of the cations (lead and
potassium) look like when they react with H2S. Be sure to make careful observations in your
notebook. Later, you will use this test to determine which cation is present in a solution by
repeating the experiment and comparing your observations.
4. To each of two clean, dry 4-mL test tubes, add 20 drops of thioacetamide and 2 drops
of nitric acid.
5. To one tube add 5 drops Pb(NO3)2, and to the other 5 drops KNO3. Heat both tubes in
a water bath for at least 5 minutes and record your observations. On heating, the thioacetamide will decompose to form H2S.
C. Identification of ions in the precipitate.
In this section of the experiment you will test the product that formed when potassium
iodide and lead nitrate solutions were mixed. To accomplish this you will re-dissolve the
solid and determine which ions it is comprised of.
Formation of Precipitate
6. Once again, add 5 drops of potassium iodide and 5 drops lead nitrate to a 4-mL test
tube. (You may simply use the sample from Part A if you still have it.) Centrifuge the
test tube, and then carefully remove and discard the clear aqueous layer. Wash the precipitate: add 1 mL of distilled water, stir, centrifuge the test tube, and discard the clear
aqueous layer. Repeat the washing.
Test for Anions
7. Add 5 drops of nitric acid, 10 drops of dichloroethane, and 5 drops of hydrogen peroxide to the solid in the test tube. Stopper and shake the test tube.
8. If necessary, alternate holding the test tube in a stream of hot tap water for 20 seconds
and shaking for 40 seconds until the yellow solid is completely dissolved. Centrifuge
the test tube. Record your observation of the dichloroethane layer and compare it with
\RXU¿QGLQJVIURP3DUW%
21
Laboratory 3
Test for Cations
9. Transfer the less-colored aqueous layer to a clean 4-mL test tube using a micropipet.
Discard the remaining dichloroethane layer in a waste beaker.
10. Heat the aqueous solution in a water bath to destroy any unreacted hydrogen peroxide.
Continue heating until the solution is completely colorless.
11. Add 20 drops of thioacetamide to the test tube, stir, and heat in a water bath. While
waiting for a reaction to occur, start on the next step. After 5 minutes, compare your
results with the reaction observed in Part B.
D. Determination of the ratio of ions in the precipitate.
In Part C, you determined which ions are in the precipitate. In this section you will make
observations and determine the ratio of these ions in the solid sample.
Preparing Test Tubes with Varying Ratios of the Reactants
/DEHO¿YHP/WHVWWXEHVWKURXJK7KH¿YHWHVWWXEHVVKRXOGEH¿OOHGFDUHIXOO\
according to the table below. Notice the table has several empty cells. Reproduce the
WDEOHLQ\RXUQRWHERRNDQG¿OOLQWKHHPSW\FHOOVbefore continuing. In each case make
the amount of lead nitrate constant and the total number of drops constant.
Table 3.1 Solutions to be mixed in Part D.
Drops
of Lead
Nitrate
Drops of
Potassium
Iodide
Drops of
Water
Lead Nitrate:
Potassium Iodide
Ratio
Tube 1
12
4
32
3:1
Tube 2
12
6
Tube 3
12
12
Tube 4
Tube 5
2:1
24
12
36
1:2
0
13. The same medicine dropper should be used to measure all drops. (Do not use a miFURSLSHW²WKHGURSVL]HVYDU\WRRJUHDWO\0HDVXUHRXWDOO¿YHDPRXQWVRIHDFKVROXtion, and then rinse the medicine dropper well with distilled water before changing
solutions.
6WRSSHUDQGVKDNHWKH¿YHWHVWWXEHV$OORZWKHSUHFLSLWDWHWRVHWWOHIRUDIHZPLQXWHV
and then centrifuge the test tubes. The centrifuge must be balanced using pairs of test
WXEHV²WU\WRFHQWULIXJHDVPDQ\WXEHVDVSRVVLEOHDWWKHVDPHWLPH
5HFRUGWKHDSSHDUDQFHDQGUHODWLYHDPRXQWVRISUHFLSLWDWHLQHDFKRIWKH¿YHWHVWWXEHV
in your notebook.
22
Development of an Equation
Determining which Test Tube, 1–5, Does Not Have an Excess of Cation or Anion
When the ions in the potassium iodide and lead nitrate solutions were mixed they may have
produced different amounts precipitate, as noted in step 4. For a given ratio, were any of the
ions in excess? If so, they will still be in the supernatant. This can be tested for by taking
a portion of the supernatant and seeing if it is still possible to form additional precipitate.
Obviously, if more precipitate is formed, excess ions must have been present in supernatant
(step 2).
16. To test a given supernatant for excess ions, use a medicine dropper and remove 5 drops
of the solution into a different test tube. Be careful not to transfer any of the precipitate.
You may need to re-centrifuge the tube to avoid transferring any solid.
•
If you are testing for the presence of potassium or iodide ions in the supernatant,
add 5 drops of lead nitrate and record the relative amount of precipitate that forms
(if any).
•
If you are testing for the presences of lead or nitrate ions in the supernatant, add 5
drops of potassium iodide.
2QFHDJDLQPDNHFDUHIXOREVHUYDWLRQV<RXPD\¿QGLWKHOSIXOWRFHQWULIXJHWKHVDPSOHV
and make comparisons.
Waste Disposal
All solutions containing lead ions and dichloroethane must be collected in a beaker at your
GHVN$GGLWWRWKHLQRUJDQLFZDVWHEHDNHULQWKHKRRGDQG¿OOLQWKHZDVWHGLVSRVDOVKHHW
Your lab instructor will neutralize the solution with sodium hydroxide and dispose of the
total volume in the appropriate container. Dispose of the used test tubes in the uncleanable
glass waste container.
23
Laboratory 3
24
Laboratory 3 Report Sheet
Development of an Equation
Name: _________________________________ Date: ______________ TA Name: ____________________
A. Initial observations.
What atoms, ions, or molecules do you think were in the potassium iodide and lead nitrate solutions
before they were mixed?
List your observations upon mixing:
Is there evidence a chemical reaction occurred? Support your answer.
Is there evidence that a gas was formed during the reaction?
B. Observations of qualitative tests for ions in solution.
Testing the Anions
Observations for the iodide reaction with H2O2:
Observations for the nitrate reaction with H2O2:
Testing the Cations
Observations for the lead reaction with H2S (thioacetamide):
Observations for the potassium reaction with H2S:
25
Laboratory 3 Report Sheet
C. Identification of ions in the precipitate that forms when lead nitrate solution is
added to potassium iodide solution.
Testing for Anions
Observations for the reaction with H2O2:
Inferences:
What anions are in the precipitate? How do you know?
Testing for Cations
Observations for the reaction with H2S:
Inferences:
What cations are in the precipitate? How do you know?
Pattern of Reactivity
What kind of reaction (combination, decomposition, combustion or exchange) occurred when potassium iodide and lead nitrate solutions are mixed? Support your reasoning.
26
Development of an Equation Report Sheet continued
D. Determining the exact ratio of ions in the precipitate.
Testing of
Precipitate
Observations on Initial Precipitate Formed (relative amount, etc.)
Tube 1
Tube 2
Tube 3
Tube 4
Tube 5
Testing of
Supernatant
Observations
Inferences: Which ions were in the
supernatant?
List both cations and anions.
Tube 1
Tube 2
Tube 3
Tube 4
Tube 5
27
Laboratory 3 Report Sheet
Using all of your observations, what is the correct ratio of ions? Support your conclusion.
Formula of the Precipitate
Write the balanced equation for the reaction that forms this precipitate.
What is your reasoning? You must convince the reader by clearly summarizing your
observations and how they lead to your inferences. This is the key point
in this lab!
28
Development of an Equation Report Sheet continued
Draw particle diagrams consistent with your balanced equation.
Before the reaction
After the reaction
Metals like lead frequently have different oxidation states, e.g., being +2 or +3,
and these oxidation states may change during a chemical reaction. What is the
oxidation state of lead in the precipitate? Support your reasoning with your
experimental observations.
29
Laboratory 3 Report Sheet
30
Laboratory 3 Answer Clinic
Development of an Equation
Question 3.1: Write a balanced chemical equation for N2 reacting with O2 to form laughing gas.
Answer
2 N2(g) + O2(g) $ 2 N2O(g)
In the boxes below, draw a particle diagram showing the reactants (before) and the products (after)
the reaction. The drawing should be consistent with your balanced equation.
Reactants (Before)
2N2(g)
+
O2(g)
Products (After)
2N2O(g)
Question 3.2: Is the chemical reaction N2(g) + O2(g) $ 2 NO(g) an example of a combination,
decomposition, combustion, or exchange reaction?
Answer
Combination reactions, which describe the reaction of two or more reactants to form one product.
Questions 3.3: What is the limiting reactant for each ratio H2:O2?
Answer
H2:O2 Ratio
Limiting Reactant
4:1
O2(g)
3:1
O2(g)
2:1
QD
1:1
H2(g)
1:2
H2(g)
31
Laboratory 3 Answer Clinic
Question 3.4: Suppose H2(g) and O2(g) are combined in the ratios shown in the table on page 17
to produce H2O. For which ratios of H2:O2 would the splint test result in a “pop.” For which ratios
of H2:O2 would the splint reignite? How does this analysis lead the chemist to the correct chemical
equation?
Answer
When the H2:O2 = 2 there will not be any H2 or O2 left over.
If H2:O2 > 2 there is excess H2 and the O2 is the limiting reactant, and the splint tests produce a
“pop.”
If H2:O2 < 2 there is excess O2 and the H2 is the limiting reactant, and the splints will be reignited.
Question 3.5: Identify the ions present in the supernatant (Ba2+, Clí, Na+DQGRU624í) prior to
additional barium chloride or sodium sulfate.
Answer
Tube 1: Ba2+ , Na+, Cl–
Tube 2: Ba2+ , Na+, Cl–
Tube 3: Na+, Cl–
Tube 4: Na+, Cl–, SO42–
Tube 5: Na+, Cl–, SO42–
Question 3.6: Use the information in the table on page 19 to determine the empirical formula of
barium sulfate. Support your reasoning.
Answer
By testing the supernatant it appears that tube 3 does not have excess barium ions or excess sulfate
ions. This tube was prepared using equal drops of barium chloride and sodium sulfate. This suggests that the ratio of barium to sulfate is 1:1, consistent with an empirical formula of BaSO4.
Question 3.7: Write a balanced chemical equation for the reaction that occurs when barium chloride and sodium sulfate are mixed.
Answer
BaCl2(aq) + Na2SO4(aq) $ BaSO4(s) + 2 NaCl(aq)
32