General Chemistry II Experiments
The Laboratory Manual for Ch12
Department of Chemistry
Ateneo de Manila University
Compiled by:
Marites J. Pasuelo
Edited by:
Edward T. Chainani
2004
To the Student:
Chemistry is an experimental science. It will be much easier to appreciate and comprehend the
science of chemistry if you actually participate in experimentation. You will find that laboratory work to be
interesting and very rewarding if you have the patience to develop good technique, the honesty to report
results as you obtain them, the initiative to deal with situations which are not covered in the text, and the
aptitude to think about the meaning of the results in the light of the chemical principles behind it.
Laboratory Safety and Behavior
The experiments contained in this manual have been thoroughly tested, and are safe when
performed properly, but there is always an element of danger (the same can be said for an automobile).
Proper behavior and proper technique in the lab will help avoid any mishap. Take heed of any special
precautions to be observed in an experiment that will be pointed out to you. Your teacher will outline the
rules of behavior that you are expected to follow, but keep in mind that all rules are based on common
sense: Do nothing that would interfere with the work of your neighbors. Clean up any spill. Report
anything not functioning properly. Leave your work area clean when you finish.
Preparation
The lab instructor is there to guide you. Don’t be afraid to ask questions. However, the lab periods
are only long enough to do the experiment: there will not be enough time if you come to the laboratory
unaware and have to figure out what you have to do. You should read the experiment procedure well in
advance so that you can get the most out of the lab. As you progress in your proficiency in the lab, fewer
directions will be given. It helps to make an outline of the procedure and plan in detail what you must do to
perform the experiment.
Laboratory Reports
Keeping a proper laboratory record is a vital aspect of good laboratory practice. Write down the
data in your lab notebook as soon as it is obtained. If you make a mistake do not obliterate your mistake,
but instead draw a simple line through it and write the correction beside it. You should always be able to
read
what
you
had
written
the
first
time.
The workbook format of this manual provides data sheets for each experiment and guides you as
to what must be recorded. The consistent tabular format is there primarily to help the teacher check
numerous (30 or more!) reports more efficiently. We hope that you, like a well-trained scientist, will not limit
yourself to putting down only the data required by the data sheet, but that you make use of blank spaces to
take note of any other information that you judge to be relevant to the experiment.
Experiment 1
CHEMICAL MASTERMIND
Objectives
To observe the changes that occur when certain ions react with one another.
To learn how to use such changes to identify ions.
Reagents
HgCl2 (mercuric chloride), Pb(Ac)2 (lead acetate), Ba(NO3)2 (barium nitrate), KI (potassium iodide) , Na2CO3 (sodium
carbonate), and Ca(NO3)2 (calcium nitrate), CuCl2 (copper (II) chloride), CuSO4 (copper (II) sulfate.
Materials
Wax paper
Procedure
Part I:
You may be familiar with the game of mastermind, in which you have to guess a color sequence of pegs within a
certain number of steps. In this experiment, you will play the game of chemical mastermind.
Here are the rules of the game:
1.
2.
3.
4.
5.
6.
7.
8.
Six solutions are labeled A, B, C, D, E, F. Each solution may contain any one of the following substances: HgCl2
(mercuric chloride), Pb(Ac)2 (lead acetate), Ba(NO3)2 (barium nitrate), KI (potassium iodide) , Na2CO3 (sodium
carbonate), and Ca(NO3)2 (calcium nitrate), but you do not know which solution contains which.
Each solution contains ions. For example, mercuric chloride contains the mercuric ion and the chloride ion. The
particular solutions have been chosen such that when any two solutions are mixed, the different ions react to form a
new product. The chemical reaction is indicated by changes, such as a change in color or the appearance of a solid or
both.
The new product may be any of the following:
HgCO3
Mercuric carbonate
Rust-colored solid
HgI2
Mercuric iodide
Orange solid
PbCO3
Lead (II) carbonate
White solid
PbCl2
Lead (II) chloride
White solid
PbI2
Lead (II) iodide
Pale yellow solid
BaCO3
Barium carbonate
White solid
CaCO3
Calcium carbonate
White solid
The object of the game is to identify the different substances A, B, C, D, E, F by mixing them together in different
combinations and observing the products that form.
Take a piece of paper of the same size as the wax paper. Draw a 6 x 6 grid and label the grid. See DATA AND RESULTS
for the pattern.
Place the wax paper over the grid you just made.
Place 1 drop each of 6 unknowns in the different squares of the grid as follows:
a)
Put 1 drop of solution A in each square of row A (horizontal).
b)
Then add 1 drop of solution B in each square of row B (horizontal).
c)
Repeat until all the horizontal rows have been "filled" with solutions.
d)
Do similarly for each vertical column.
Record your observations.
CHEMICAL MASTERMIND
-4-
9.
10.
To differentiate between Ca(NO3)2 and Ba(NO3)2, add 1 drop of 3M H2SO4. BaSO4 is a white precipitate that forms from
the reaction of Ba(NO3)2 and H2SO4.
Ca2+ ions do not form a precipitate with H2SO4.
Part II
The procedure is the same as above but we will be using the following solutions: potassium chloride, KCl; silver
nitrate, AgNO3; barium nitrate, Ba(NO3)2; copper (II) chloride, CuCl2; and copper (II) sulfate, CuSO4. In addition, you need to
know that silver chloride; silver sulfate and barium sulfate are white precipitates. You should also be aware that copper salts in
solution give the solution a clear blue color.
Questions:
1.
2.
3.
Think about this experiment as a whole. Which observations did you find most useful in leading you to your conclusion?
Try to think back on this experiment and recall any observation that you may have made but did not think it important
enough to write down. What are these, if any?
If there were observations you made but did not record, how did that affect your work?
1
CHEMICAL MASTERMIND
Data Sheet
Part I: Observations
A
B
C
A
B
C
D
E
F
Identification of Unknown solutions:
A / ___ =
___________________________________
B / ___ =
___________________________________
C / ___ =
___________________________________
D / ___ =
___________________________________
E / ___ =
___________________________________
F / ___ =
___________________________________
D
E
F
Experiment 2
RATE OF REACTIONS: EFFECT OF
CONCENTRATION AND TEMPERATURE
Objectives
To find out how temperature and concentration affect the rate of reaction.
Reagents
0.1 M sodium thiosulfate (‘hypo’), distilled water, 2M hydrochloric acid.
Materials
Water bath, test tubes, 10 mL graduated cylinder, thermometer, watch with readout in seconds. (Note: students
should bring their own watch.)
Procedure
This experiment has two parts. In the first part, we will study the effect of concentration on the rate of reaction and on
the second part, the effect of temperature on the rate of reaction. In both parts, we shall make use of the reaction of sodium
thiosulfate with hydrochloric acid (See equation below).
3 Na2S2O3(aq)
+
2 HCl(aq) 
4 S(s)
+
other products
When the thiosulfate and hydrochloric acid are added together, the time it takes for sulfur to form can be taken as an indication of
the rate of reaction of the two.
In the first part, if successive runs of the same amount of acid but different amounts of sodium thiosulfate solution are
used, the time it takes for sulfur to form in the different runs gives us information on how the concentration affects the speed of
reaction. In the second part, we are going to study the rate of reaction at a fixed concentration of reactants but at different
temperatures.
Part I (Effect of concentration)
1. Label the 5 test tubes A, B, C, D, E.
2. Prepare sodium thiosulfate solution of different concentrations by mixing “hypo” solution and distilled water according to the
table below.
Solution
A
B
C
D
E
3.
4.
5.
6.
0.1 M Na2S2O3
5 mL
4 mL
3 mL
2 mL
1 mL
Distilled water
0 mL
1 mL
2 mL
3 mL
4 mL
To each solution, in turn, quickly add 5 mL of 2M HCl and immediately give it one hard vertical shake. Take as time zero
the time of the shake.
Watch for the first appearance of cloudiness. (Cloudiness is best observed if the test tube is held against a dark
background.) Note the time (in seconds) when cloudiness first appears.
Do three trials for each concentration of thiosulfate. Clean the test tubes thoroughly with a test tube brush before using
them again.
Calculate the average time, t and average 1/t.
Graph concentration of sodium thiosulfate vs. time and concentration of sodium thiosulfate vs. 1/t.
Part II (Effect of Temperature)
1. Prepare a stock solution of Na2S2O3 by measuring 30mL of 0.1M Na2S2O3 and adding to it 20 mL distilled water. Mix well.
Use this solution for the following experiment. (Note: If 0.06 M Na2S2O3 is available, you don’t have to prepare the stock
solution, use the 0.06M directly.)
RATE OF REACTION: EFFECT OF CONCENTRATION AND TEMPERATURE
-7-
2.
3.
4.
5.
6.
Do the experiment at the following temperatures:
Run 1:
room temperature (rT)
Run 2:
200C lower than room temperature
Run 3:
100C lower than room temperature
Run 4:
100C higher than room temperature
Run 5:
200C higher than room temperature
Note: appropriate mixing of ice-cold or hot water with tap water can prepare water baths at the desired temperatures.
Perform each run twice.
a) Place 5 mL Na2S2O3 stock solution in test tube A.
b) Measure 5 mL 2M HCl into test tube B.
c) Place both test tubes A and B in a water bath set at the desired temperature until the temperature of the solutions in
the test tubes have reached the temperature of the bath (about 5-10 minutes).
d) Quickly add the hydrochloric acid (test tube B) to the sodium thiosulfate solution (test tube A). Shake once quickly and
note the time. Determine the time when the first sign of cloudiness appears.
e) Repeat at the other temperature.
Calculate the average time it takes for the reaction to occur and the average 1/t.
Graph the reciprocal of average time, 1/t vs. temperature.
Also graph ln (1/t) vs. 1/temp. ln is the natural logarithm (logarithm to the base e).
Questions
1. From the graph, what is the relationship between concentration and time?
2. From the graph, what is the relationship between concentration and the rate of the reaction, 1/t. Is your answer here
consistent with your answer above? Explain.
3. As a result of this experiment, can you say
a. the rate depends on concentration, or
b. concentration depends on rate, or
c. both of the above, or
d. neither of the above
Explain.
4. From your data, what is the relationship, between temperature and the rate of reaction? Does this support or contradict the
Kinetic Molecular Theory? Explain your answer.
5. Using your graph in temperature vs. 1/t, estimate the change in temperature needed to double the reaction rate; to decrease
by half. Is the change in temperature a constant value? What is its significance?
2
RATE OF REACTION:
CONCENTRATION & TEMPERATURE
Data Sheet
Name: _____________________________
Date: ______________________
Part 1
Solution
Trial 1
Time (s)
Trial 2
Average time, t
Trial 3
Average 1/t
A
B
C
D
E
Part 2
Run
1
2
3
4
5
Time (s)
Temperature, T
Trial 1
rT
rT - 200C
rT - 100C
rT + 100C
rT + 200C
Trial 2
Average time, t
(s)
Average 1/t (1/s)
Experiment 3
TITRATION OF ACIDS AND BASES
Objectives
To become familiar with the techniques of titration, a volumetric method of analysis. To determine the amount of acid in
an unknown by titration with a base.
Reagents
0.50 M NaOH, phenolphthalein indicator, unknown acid
Materials
50 mL buret, buret clamp, ring stand, wash bottle, 250 mL Erlenmeyer flasks
Procedure
One of the most common and familiar reactions in chemistry is the reaction of an acid with a base. This reaction is
termed neutralization, and the essential feature of this process in aqueous solution is the combination of hydronium ions with
hydroxide ions to form water:
H30+(aq) + OH-(aq)  2H2O(l)
In this experiment you will use this reaction to determine accurately the concentration of a hydrochloric acid solution.
You will do this by reacting it with a base of accurately known concentration until it is neutralized. The stockroom personnel have
accurately determined or standardized, the concentration of the base that you will use (see Note 1). Thus, all that needs to be
done is to measure the amount of acid in an unknown. To do this, you will accurately measure with a buret the volume of
standard base that is required to exactly neutralize the acid present in the unknown. The technique of accurately measuring the
volume of a solution required to react with another reagent is termed titration.
An indicator solution is used to determine when an acid has exactly neutralized a base, or vice versa. A suitable
indicator changes colors when equivalent amounts of acid and base are present. The color change is termed the end point of
the titration. Indicators change colors at different pH values. Phenolphthalein, for example, changes color from colorless to pink
at a pH of about 9; in slightly more acidic solutions it is colorless, whereas, in more alkaline solutions it is pink. The point at which
stoichiometrically equivalent quantities are brought together is known as the equivalence point of the titration.
It should be noted that the equivalence point in a titration is a theoretical point. It can be estimated by observing some
physical change associated with the condition of equivalence, such as the change in color of an indicator, which is termed the
end point.
The most common way of quantifying concentrations is molarity (symbol M), which is defined as the number of moles of
solute per liter of solution, or the number of millimoles of solute per milliliter of solution:
M
=
Moles solute
--------------------------------------Volume of solution in liters
=
10-3 mole
--------------10-3 liter
=
mmol
--------mL
[1]
From above, you can solve for the moles of solute in any volume (in liters) of known molarity as follows:
moles solute = M solution x V solution
[2]
TITRATION OF ACIDS AND BASES
- 10 -
To find the number of moles of base (NaOH) in a volume of base solution, Vb, of known molarity, Mb, use the following
equation:
moles base = Mb x Vb
[3]
If the acid is HCl, one mole NaOH will neutralize one mole of acid. Hence, the moles of NaOH used to neutralize a given
volume of HCl will be equal to the number of moles of acid in that volume:
moles base = moles acid
[4]
To find the molarity of the acid:
Mb x Vb = Ma x Va
Ma = (Mb x Vb) / Va
[5]
[6]
Preparation of a Buret for Use
The 50-mL buret is supplied to you clean and ready to use. Normally it is stored with water to maintain a clean inside
surface. Discard this water and rinse the buret with at least five 10-mL portions of distilled water. The water must run freely
from the buret without leaving any drops adhering to the sides. Make sure that the buret does not leak and that the
stopcock turns freely.
Reading a Buret
All liquids, when placed in a buret, form a curved meniscus at their upper surfaces. In the case of water or water solutions,
this meniscus is concave, and the most accurate buret readings are obtained by observing the position of the lowest point
on the meniscus on the graduated scales.
To avoid parallax errors when taking readings, the eye must be on a level with the meniscus. Wrap a strip of paper around
the buret and hold the top edges of the strip evenly together. Adjust the strip so that the front and back edges are in line
with the lowest part of the meniscus and take the reading by estimating to the nearest tenth of a marked division (0.01 mL).
Analysis of an Unknown Acid
Prepare about 100-150 mL of C02-free water by boiling for about 5 minutes (see Note 2). When it cools, transfer it into
your wash bottle. You can use this for washing down drops of your titrant (NaOH) during titration.
Rinse the previously cleaned buret with at least four 5-mL portions of the approximately 0.500 M sodium hydroxide
solution that you have prepared. Discard each portion. Do not return any of the washings to the bottle. Completely fill the buret
with the solution and remove the air from the tip by running out some of the liquid into an empty beaker. Make sure that the
lower part of the meniscus is at the zero mark or slightly lower. Allow the buret to stand for at least 30 seconds before reading
the exact position of the meniscus. Remove any hanging drop from the buret tip by touching it to the side of the beaker used for
the washings. Record the initial buret reading.
Place each unknown in the wide-mouthed 250 mL Erlenmeyer and add two drops of phenolphthalein indicator solution.
Titrate with your standard sodium hydroxide solution to the faintest visible shade of pink (not red) in the following manner:
As the sodium hydroxide solution is added, a pink color appears where the drops of the base come in contact with the
solution. This coloration disappears with swirling. As the end point is approached, the color disappears more slowly, at which
time the sodium hydroxide should be added drop by drop. It is most important that the flask be swirled constantly throughout the
entire titration. The end point is reached when one drop of the sodium hydroxide solution turns the entire solution in the flask
from colorless to pink. The solution should remain pink when it is swirled. Allow the titrated solution to stand for at least I minute
so the buret will drain properly. Remove any hanging drop from the buret tip by touching it to the side of the flask and wash
down the sides of the flask with a stream of water from the wash bottle. Record the buret reading. Repeat this procedure with
the other two samples. Do at least three runs. For good results the determinations should agree within 1.0 percent. Your
answers should have four significant figures. Compute the standard deviation of your results.
TITRATION OF ACIDS AND BASES
- 11 -
Test your results by computing the average deviation from the mean. If one result is noticeably different from the
others, perform an additional titration. If any result is more than two standard deviations away from the mean, discard it and
titrate another sample.
Questions
1. What are equivalence points and end points and how do they differ?
2. What is parallax and why should you avoid it?
3. What is the effect on your result of overshooting the endpoint?
Notes
1.
In this experiment your solution of NaOH was standardized by titrating it against a very pure sample of potassium
hydrogen phthalate, KHC8H404, of known weight. Potassium hydrogen phthalate (often abbreviated as KHP) has one acidic
hydrogen. Its structure is shown below. It is a monoprotic acid with the acidic hydrogen bonded to oxygen and has a molar
mass of 204.2 g.
OK
O
O OH
The balanced equation for the neutralization of potassium hydrogen phthalate is given below:
KHC8H404(aq)
+ NaOH(aq)

H2O(l) + KNaC8H404(aq)
[N-1]
In the titration of the base NaOH against KHP, an equal number of moles of KHP and NAOH are present at the equivalence
point, In other words, at the equivalence point
Moles NaOH = Moles KHP
[N-2]
Because primary standard KHP, a very pure substance, is too costly to be used by every student, the trained stockroom
personnel carry out the standardization on the entire batch of NaOH that the class uses.
2.
Carbon dioxide from the atmosphere dissolves in water to form carbonic acid. This weak acid reacts with the titrant
NaOH and thus causes an error during titration. Carbon dioxide may be removed from water by boiling it for 5 minutes.
3
TITRATION OF ACIDS AND BASES
Data Sheet
Name: _____________________________
Date: _____________________
Analysis of an Unknown Acid
Trial 1
Molarity of base (from instructor) (M)
Volume of the unknown acid (mL)
Initial volume reading for NaOH (mL)
Final volume reading for NaOH (mL)
Total volume of NaOH delivered (mL)
Molarity of the Unknown Acid (M)
Average Molarity of the Unk. Acid: _________________
Standard deviation of unknown: __________________
Show your calculations of Molarity and standard deviation:
Trial 2
Trial 3
Trial 4
Experiment 4
PREPARATION OF BUFFERS
Objectives
To become familiar with buffer systems and buffer calculations.
Required Reading
Please read the pH meter user’s manual (available from the stockroom) before the lab session.
Introduction
A buffer is a solution consisting of a definite proportion of conjugate base [A-] to its weak acid [HA] with a pH near the
pKa of the weak acid. Since the ratio [A-]/[HA] ranges from 10-1 to 101, the two species are always present in considerable
amounts. Together they resist large change in pH by partially neutralizing the H+ and OH- ions added to the system, as shown
by the following equations:
H+ + AOH- + HA 

HA
A- + H2O
(1)
(2)
In the
HendersonHasselbach’s equation (eqn. 3), the most effective buffering system contains equal concentration of [HA] and [A-]. When [A-] is
equal to [HA], pH equals pKa (eqn. 4). The effective range of a buffer system is generally two pH units, centered at the pKa
value.
pH = pKa + log [A-]/[HA]
(3)
when [A-]=[HA], pH = pKa + log 1
pH = pKa
(4)
Reagents
3M acetic acid, pH meter calibration buffers (pH =4, 7 and 10), sodium acetate, dilute HCl, dilute NaOH, pipets
Equipment
pH meter, three 50-mL volumetric flasks, wash bottle, beakers, balance, pipets
Procedure:
This experiment will be performed in groups of 5 or 6. The group will have the same raw data but data processing and
analysis of results must be done individually. Show all calculations in each part.
I. Calibrating the pH meter
Read the short notes on how to use the pH meter and the pH meter manual (usually beside the pH meter logbook).
Calibrate the pH meter using three buffers (pH=4, 7, and 10). Make sure the buffers are warmed up to room
temperature.
II. Determining the experimental Ka
In this part, one must prepare a 50-mL, 0.12M acetic acid buffer from a 3M acetic acid stock solution and solid
sodium acetate.
First, calculate how much 3 M acetic acid stock is needed. Check the available volumetric pipets in the laboratory and
determine which one can be used to obtain the desired volume. Pour a small amount of the acetic acid stock from the stock
PREPARATION OF BUFFERS
- 14 -
solution bottle into a clean, dry beaker. Pipet out the desired amount and put it in another beaker. Never pipet directly from the
stock solution bottle.
Based on the desired final buffer concentration, calculate the mass of sodium acetate that is needed such that the
concentrations of both acetic acid and sodium acetate are equal. Weigh out the appropriate amount of sodium acetate and
transfer it to the beaker with the pipeted acetic acid stock. Add distilled water, if necessary, until the solid completely dissolves.
Transfer the solution to a 50-mL volumetric flask. Add distilled water to the 50-mL mark and mix well.
Measure the pH of the buffer. From the pH, determine the experimental Ka for acetic acid. Note down the room
temperature as well.
III. Preparing buffers of a specific pH
In part II, the ratio of salt to acid was 1:1. For this part, one must prepare a 50-mL, 0.12M acetic acid buffer of a
specific pH assigned to the group from a 3M acetic acid stock solution and solid sodium acetate.
Using the Ka obtained from part II, calculate the mass of sodium acetate that will be added to the acetic acid stock to
arrive at the assigned pH. Make a 50.0-mL sample of the buffer using the same procedure as in part II.
Check the pH of the buffer.
IV. Determining the effect of a strong acid or base to buffers
Transfer half of the buffer from part III into a clean, dry beaker. Note down its exact volume. On one part of the buffer,
add 1 mL of dilute HCl. Use a dry stirring rod to thoroughly mix the acid and the buffer. Check the pH. Add 1-mL increments until
you have added a total of 5 mL acid. Measure the pH at each stage. Do the same for the base using the other half of the buffer.
Results and Discussion
What is the Ka of acetic acid at 298 K? Compare it with the Ka obtained in part II and account for the difference.
In part III, were you able to obtain the desired pH? Account for the difference.
Calculate the expected pH after adding 1-mL increments of the acid and the base in part IV. Compare your results to
the measured pH at each stage.
Questions:
1. Would a solution made by mixing HCl and NaOH be an effective buffer? Explain.a
2. In part IV, adding a strong acid increases the amount of weak acid in solution and decreases that of the salt. Conversely,
adding a strong base decreases the amount of the weak acid and increases that of the salt. How can you claim that the buffered
solution resists changes in the pH? a
References:
Boyer, R. 1993. Modern Experimental Biochemistry, 2nd ed. The Benjamine/Cummings Publishing Company, Inc.
http://www4.cord.edu/chemistry/pmork/chem137/DesigningBuffer.doc, last accessed 2003
a
Zumdahl S. 1998. Chemical Principles, 3/e. Houghton Mifflin Company: pp. 331.
Experiment 5
CHEMISTRY OF GROUP 1 CATIONS
Introduction
Because the chlorides of Pb2+, Ag+, and Hg22+ are insoluble, they may be precipitated and separated from the cations
of groups 2, 3 and 4 by the addition of HCl. The following equations represent the reactions that occur:
Pb2+(aq) + 2 Cl- (aq) 
Ag+(aq) + Cl- (aq)

Hg22+(aq) + 2 Cl- (aq) 
PbCl2 (s)
AgCl (s)
Hg2Cl2 (s)
white
white
white
[1]
[2]
[3]
A slight excess of HCl is used to ensure complete precipitation of the cations and to reduce the solubility of the chlorides by the
common-ion effect. However, a large excess of chloride must be avoided, because both AgCl and PbCl2 tend to dissolve by
forming soluble complex anions:
PbCl2 (s) + 2Cl-(aq)
AgCl(s) + Cl-(aq)


PbCl42-(aq
AgCl2-(aq)
[4]
[5]
PbCl2 is appreciably more soluble that either AgCl or Hg2Cl2. Thus, even when PbCl2 precipitates, a significant amount of Pb2+
remains in the solution and is subsequently precipitated with the group 2 cations as the sulfide PbS. Because of its solubility,
Pb2+ sometimes does not precipitate as the chloride, because either its concentration is too small or the solution is too warm.
Lead. Lead chloride is much more soluble in hot water than in cold. It is separated from the other two insoluble chlorides by
dissolving it in hot water. The presence of Pb2+ is confirmed by the formation of a yellow, PbCrO4, upon the addition of K2CrO4:
Pb2+ + CrO42-(aq)

PbCrO4(s)
yellow
[6]
Mercury(I). Silver chloride is separated from Hg2Cl2 by the addition of aqueous NH3. Silver chloride dissolves because Ag+
forms a soluble complex cation with NH3:
AgCl(s) + 2NH3(aq)

Ag(NH3)2+(aq) + Cl-9aq)
[7]
Mercury(I) chloride reacts with aqueous ammonia in a disproportionation reaction to form dark gray precipitate:
Hg2Cl2(a) + 2NH3(aq)  HgNH2Cl(s) + Hg(l) + NH4+(aq) + Cl-(aq)
[8]
Although HgNH2Cl is white, the precipitate appears dark gray because of the colloidal dispersion of Hg(l).
Silver. To verify the presence of Ag+, the supernatant liquid from the last reaction is acidified and AgCl reprecipitates if Ag+ is
present. The acid decomposes Ag(NH3)2+ by neutralizing NH3 to form NH4+. It is necessary that the solution be acidic, or else
the AgCl will not precipitate and Ag+ can be missed.
Ag(NH3)2+(aq) + 2H+(aq) + Cl-(aq)  AgCl(s) + 2NH4+(aq)
[9]
Procedure:
First you will analyze a known that contains all three cations of group 1. Record on your report sheet the reagents
used in each step, your observations, and the equation for each precipitation reaction. After completing the analysis of a known,
obtain an unknown. Follow the same procedures as with the known. Also record conclusions regarding the presence or
absence of all cations. Before beginning this experiment review the techniques used in qualitative analysis, eg., centrifugation,
heating solutions, washing precipitates, and testing acidity.
CHEMISTRY OF GROUP 1 CATIONS
G1-1.
- 16 -
Precipitation of Group 1 Cations
Measure out 10 drops (0.5 mL) of the test solution or the unknown into a small (10 mm x 75 mm) test tube. Add 4 drops of 6M
HCl, stir thoroughly, and then centrifuge. Test for completeness of precipitation by adding 1 drop of 6M HCl to the clear
supernate. If the supernate turns cloudy, this shows that not all of the group1 cations have precipitated; add another 2 drops of 6
M HCl, stir and centrifuge. Repeat this process until no more precipitate forms. All of the group 1 cations must be precipitated or
else they will slip through and interfere with subsequent group analyses. If a general unknown is being analyzed, decant the
supernate into a clean test tube and save it for analysis of group 2 cations; otherwise, discard it. Wash the precipitate by adding
5 drops of cold distilled water and stirring. Centrifuge and add the liquid to the supernate.
G1-2.
Separation and Identification of Pb2+
Add 15 drops of distilled water to the precipitate and place the test tube in a hot water bath. Stir using a stirring rod and heat for
1 min. or longer. Quickly centrifuge and decant the hot supernate into a clean test tube. Repeat this procedure two more times,
combining the supernates, which should contain Pb2+ if it is present. Save the precipitate for procedure G1-3. Add 3 drops of
1M K2CrO4 to the supernate. The formation of a yellow precipitate, PbCrO4, confirms the presence of Pb2+.
G1-3.
Separation and Identification of Ag+ and Hg22+
Add 10 drops of 6M NH3 to the precipitate from step G1-2. The formation of a dark gray precipitate indicates the presence of
mercury. Centrifuge and decant the clear supernate into a clean test tube. Add 20 drops of 6 M HNO3 to the decantate. Stir the
solution and test its acidity. Continue to add HNO3 dropwise until the solution is acidic. A white cloudiness confirms the
presence of Ag+.
Review Questions
1.
2.
3.
4.
What are the symbols and charge of group 1 cations?
Which chloride salt is insoluble in cold water but soluble in hot water?
Which chloride salt dissolves in aqueous NH3?
How could you distinguish;
a. BaCl2 from AgCl?
b. HNO3 from HCl?
5. Complete and balance the following equations:
a. AgCl(s) + NH3(aq) 
b. Pb2+(aq) + CrO42-(aq) 
c. Hg2Cl2 (s) + NH3 (aq) 
d. Ag(NH3)2+(aq) + H+(aq) + Cl-(aq) 
6. What can you conclude if no precipitate forms when HCl is added to the unknown solution?
7. Why are precipitates washed?
8. How do you decant supernatant liquids from small test tubes?
CHEMISTRY OF GROUP 1 CATIONS
Group 1 flow scheme. Procedure numbers are given in circles.
- 17 -
5
CHEMISTRY OF GROUP 1 CATIONS
Data Sheet
Name: _____________________________
Date: ______________________
Part I: Known
Procedure
Reagent/s
Observations
Equations
Teacher’s
Sig.
Observations
Equations
+/-
Part II: Unknown
Procedure
Reagent/s
Ions Present: ___________________________________________________
Ions Absent: __________________________________________________________
Experiment 6
CHEMISTRY OF GROUP 2 CATIONS
Introduction
Hydrogen sulfide is the precipitating agent for the group 2 cations, Pb2+, Cu2+, Bi3+, and Sn4+. We will generate H2S
from the hydrolysis of thioacetamide, CH3CSNH2:
CH3CSNH2(aq)+ 2 H2O(l)  H2S(aq) + CH3CO2-(aq) + NH4+(aq)
[1]
By controlling the hydrogen-ion concentration of the solution, we can control the sulfide-ion concentration. We see from
Equation [2], which represents the overall ionization of H2S, that the equilibrium will shift to the left if the hydrogen-ion
concentration is increased by adding some strong acid to the solution.

H2S(aq)
2H+(aq)
+
S2-(aq)
[2]
Under these conditions the sulfide-ion concentration is very small. Thus by adjusting the pH of the solution to about 0.5, only the
more insoluble sulfides will precipitate. These sulfides are those of group 2, PbS, CuS, Bi2S3, and SnS2.
Pb2+(aq)
Cu2+(aq)
2Bi3+(aq)
Sn2+(aq)
+ S2-(aq)
+ S2-(aq)
+ 3S2-(aq)
+ 2S2-(aq)




PbS(s)
CuS(s)
Bi2S3(s)
SnS2(s)
black
black
brown
yellow
[3]
[4]
[5]
[6]
The sulfides of group 3 are soluble in such acidic solutions and therefore do not precipitate because the sulfide ion concentration
is too small. The flow chart in Figure 2 summarizes how the cations of group 2 are separated and identified. Note the colors of
the precipitates.
The cations are first treated with hydrogen peroxide (H2O2) and HCl to ensure that tin is in +4 oxidation state:
Sn2+(aq) + 2H2+(aq) + H2O2(aq)  Sn4+(aq) + 2H2O(l)
[7]
Tin must be in the +4 oxidation state to form the soluble SnS32- ion and thereby allow it to be separated from the insoluble
sulfides PbS, CuS, and Bi2S3:
SnS2(s)
+
S2-(aq)

SnS32-(aq)
[8]
These insoluble sulfides are brought into solution by treating them with hot nitric acid. Elemental sulfur is formed in the oxidation
reactions:
3PbS(s) + 8H+(aq) + 2NO3-(aq)  3Pb2+(aq) + 3S(s) + 2NO(g) + 4H2O(l)
3CuS(s) + 8H+(aq) + 2NO3-(aq)  3Cu2+(aq) + 3S(s) + 2NO(g) + 4H2O(l)
Bi2S3 + 8H+(aq) + 2NO3-(aq)  2Bi2+(aq) + 3S(s) + 2NO(g) + 4H2O(l)
[9]
[10]
[11]
Lead. After the sulfides of lead, copper, and bismuth are brought into solution, lead is precipitated as the white sulfate by the
addition of H2SO4:
Pb2+(aq)
+
SO42-(aq) 
PbSO4(s)
white
[12]
The solution is heated strongly to drive off HNO3, because PbSO4 is soluble in the presence of nitric acid. To confirm the
presence of Pb2+, the PbSO4 is dissolved in NH4C2H3O2 and precipitated as yellow PbCrO4:

Pb(C2H3O2 )42-(aq) + SO42-(aq)
[13]
PbSO4(s) + 4 C2H3O2-(aq)
Pb(C2H3O2 )42-(aq) + CrO42- (aq) 
4 C2H3O2-(aq) + PbCrO4(s)
yellow
[14]
CHEMISTRY OF GROUP 2 CATIONS
- 20 -
Copper. The flow chart for Group 2 shows that addition of aqueous NH3 to the solution containing Cu2+ and Bi3+, precipitates
Bi3+ as white Bi(OH)3 and forms a soluble deep blue amine complex of copper, Cu(NH3)42+. The deep blue color confirms the
presence of copper and can be seen even when the concentration of copper is very low.
Bi3+(aq) + 3NH3(aq) + 3 H2O(l)  3NH4+(aq) + Bi(OH)3(s)
Cu2+(aq) + 4NH3(aq) + 3 H2O(l)  Cu(NH3)42+ (aq)
white
deep blue
[15]
[16]
Bismuth. The Bi(OH)3 precipitate is often difficult to observe if the solution is blue. Bismuth is confirmed by separating the
Bi(OH)3 from the solution and reducing it with SnCl2 in an alkaline solution. A black powder of finely divided bismuth is formed:
2Bi(OH)3(s) + 3Sn(OH)4 2-(s)  Sn(OH)6 2- + 2Bi(s)
black
[17]
Tin. Addition of concentrated HCl to a solution of SnS32- first yields the precipitate SnS2, which eventually dissolves the acid
solution:
SnS32-(aq) + 2H+(aq)  SnS2(aq) + H2S(g)
SnS2(s) + 4H+(aq) + 6Cl-  SnCl62-(aq) + 2H2S(g)
[18]
[19]
Aluminum metal reduces tin from +4 to the +2 oxidation state. The Sn2+ (in the form of the complex SnCl42-) in turn reduces
HgCl2 to insoluble Hg2Cl2, which is white, and Hg metal, which is black. Thus, the precipitate appears white to gray in color.
3SnCl62-(aq) + 2Al  3SnCl42-(aq) + 2AlCl3
2SnCl42-(aq) + 3HgCl2(aq)  2SnCl62-(aq) + Hg2Cl2(s) + Hg(l)
[20]
[21]
Procedure
G2-1. Oxidation of Sn2+ and Precipitation of Group 2
Place the decantate from Procedure G1-1 or 7 drops of "known" or "unknown" solution in a caserole. Add 4 drops of 3% H 2O2
and 4 drops of 2M HCl. Carefully boil the solution by passing it back and forth over the flame of your burner. Formation of brown
areas on the bottom of the casserole indicates overheating. If brown areas appear, swish the solution around until the brown
area disappears. When the volume of the solution is reduced to about 5 drops, stop heating and allow the heat from the
casserole to complete the evaporation. About 3 drops of solution should be present after cooling. Add 10 drops of 6M HCl. IN
THE HOOD evaporate the contents of the casserole to a pasty mass, being careful again to avoid overheating. Let the
casserole cool, and then add 5 drops of 2M HCl and 5 drops of distilled water. Swish the contents to dissolve or suspend the
residue and transfer it to a small test tube. [The solution is evaporated to dryness or to a pasty mass in order to remove the
unknown quantity of acid that is present. A known amount of HCl is then added that is required for the group precipitation as
sulfides.]
Add 10 drops of 1M thioacetamide to the solution in the test tube. Stir the mixture and heat the test tube in a boiling-water bath
for 10 min. If excessive frothing occurs, temporarily remove the test tube from the bath. Occasionally stir the solution while it is
being heated. After heating for 10 min, add 10 drops of hot water and 10 drops of 1M thioacetamide and 2 drops of 1M
NH4C2H3O2 (ammonium acetate). Mix and heat in the boiling-water bath for 10 more minutes, stirring occasionally. Cool,
centrifuge, and using a pipet, decant into a test tube. The precipitate contains the group 2 insoluble sulfides, while the
supernatant liquid may contain cations from groups 3 and 4.
Test the decantate for completeness of precipitation by adding 3 drops of H2O, 1 drop of 1M NH4C2H3O2, and 2 drops of 1M
thioacetamide. Mix and heat in the boiling-water bath for 1 min. If colored precipitate forms, continue heating for 3 additional
minutes. A faint cloudiness may develop because of the formation of colloidal sulfur. Repeat the precipitation procedure
precipitation is complete. If the decantate is to be analyzed for groups 3 and 4, transfer it to a casserole and boil it to reduce the
volume to about 0.5 mL. Otherwise, discard it. Transfer the contents to a labeled test tube and save it for procedure G3-1.
Rinse the casserole with 6 drops of water and add the washing to the test tube and stopper it.
CHEMISTRY OF GROUP 2 CATIONS
- 21 -
All precipitates should be combined into the test tube containing the original precipitate using a few drops of water to aid in the
transfer. Wash the precipitate three times, once with 10 drops of hot water and twice with 20-drop portions of a hot solution
prepared using equal volumes of water and 1M NH4C2H3O2. Be certain to stir the the washing liquid and precipitate with a
stirring rod before each centrifugation. Should the precipitate form a colloidal suspension, add 10 drops of 1M NH4C2H3O2 and
heat the suspension in the boiling-water bath. Discard the washings. Note, NH4C2H3O2 is often used in the washing of
precipitates. Its purpose is to help prevent the precipitate from becoming colloidal. Finely divided particles are difficult to settle.
G2-2. Separation of Sn4+ from PbS, CuS and Bi2S3
To the precipiatate for Procedure G2-1 add 10 drops of (NH4)2S (ammonium sulfide) solution and stir well. Then heat for 3 to 4
mins in the boiling water bath. Remove the test tube from the water bath as necessary to avoid excessive frothing. Centrifuge,
decant, and save the decantate. Repeat the treatment using 7 drops of (NH4)2S. Centrifuge and combine the decantate with the
first. Stopper the combined decantate and save for analysis for tin (Procedure G2-5).
Wash the precipitate twice with 20-drop portions of a hot solution prepared by mixing equal volumes of water and 1M
NH4C2H3O2. The precipitate may contain PbS, CuS and Bi2S3 and should be analyzed according to Procedure G2-3.
G2-3. Separation and Identification of Pb2+
Add 1 mL (20 drops) of 3M HNO3 to the test tube containing the precipitate from Procedure G2-2. Mix thoroughly and transfer
the contents to a casserole. Boil the mixture gently for 1 min. Add more HNO3, if necessary, to keep the amount of liquid
constant. Cool, centrifuge, and discard any free sulfur that forms. Transfer the decantate to a caserrole and add 6 drops of 18M
H2SO4. Be careful: Concentrated H2SO4 causes severe burns. If you get any on yourself, immediately wash the area with
copious amounts of water. IN A HOOD, evaporate the contents until the volume is about 1 drop and a dense white fumes of SO3
are formed. The fumes should be so dense that the bottom of the casserole cannot be seen. The appearnce of dense white
fumes of SO3 ensures that all HNO3 has been removed. Cool, add 20 drops of water, and stir. Quickly transfer the contents to a
test tube before the suspended material settles. Cool the test tube. A finely divided white precipitate, PbSO4, indicates the
presence of lead. Centrifuge and save the decantate for Procedure G2-4. Wash the precipitate twice with 10-drop portions of
cold water and discard the washings. Add 6 drops of 1M NH4C2H3O2 to the precipitate and stir for about 15s. Then add 1 drop
of 1M potassium chromate, K2CrO4. The formation of a yellow precipitate (PbCrO4) confirms the presence of lead.
G2-4. Separation of Bi3+ and Identification of Bi3+ and Cu2+
To the decantate from Procedure G2-3 carefully add 15M aqueous NH3 (avoid inhalation or skin contact) dropwise while
constantly stirring until the solution is basic to litmus. The appearance of a deep blue color of Cu(NH3)42+ confirms the presence
of Cu2+. Centrifuge and discard the supernatant liquid, but save the white precipitate of Bi(OH)3. Be careful in observing the
white gelatinous Bi(OH)3, for it may be somewhat difficult to see when the solution is colored.
Wash the precipitate once with 10 drops of hot water and discard the washings. Add 6 drops of 6M NaOH and 4 drops of freshly
prepared* 0.2M SnCl2 to the precipitate and stir. The formation of a jet-black precipitate confirms the presence of Bi3+.
G2-5. Identification of Sn4+
Transfer the decantate from procedure G2-2 to a casserole and boil for 1 min to expel H2S; then add 4 drops of cold water. Add a
1-in piece of 26 gauge aluminum wire (or a 1-inch strip of aluminum foil) and heat gently until the wire or foil has dissolved.
Continue to gently heat the solution for about 2 more minutes, replenishing the solution with 6M HCl if necessary. There should
be no dark residue at this stage; if there is, continue heating until it dissolves. (You cannot stop at this point.) Transfer the
solution to a test tube and cool under running water. Immediately add 3 drops of 0.1M mercuric chloride, HgCl2, and mix. Allow
the mixture to stand for 1 min. The formation of a white or gray precipitate confirms the presence of tin.
CHEMISTRY OF GROUP 2 CATIONS
Review Questions
1.
2.
3.
4.
5.
6.
7.
What are the symbols and charges of group 2 cations?
How could CuS be separated from SnS2?
How are Cu2+ and Bi3+ separated?
How can Ag+ be separated from Cu2+?
Complete and balance the following equations:
a. Bi3+(aq) + S2-(aq) 
b. SnS2(s) + S2-(aq) 
c. PbS(s) + H+ (aq) + NO3- 
d. Bi3+(aq) + + NH3(aq) + H2O(l) 
Why is H2O2 added in the initial step of the separation of Group 2 cations?
What is the color of CuS? Of SnS2? Of PbSO4?
- 22 -
Experiment 6
CHEMISTRY OF GROUP 2 CATIONS
Group 2 Flow scheme.
6
CHEMISTRY OF GROUP 2 CATIONS
Data Sheet
Name: _____________________________
Date: ______________________
Part I: Known
Procedure
Reagent/s
Observations
Equations
Teacher’s
Sig.
CHEMISTYR OF GROUP 2 CATIONS
- 25 -
Part II: Unknown
Procedure
Reagent/s
Observations
Equations
Ions Present: ___________________________________________________
Ions Absent: __________________________________________________________
+/-
Experiment 7
CHEMISTRY OF GROUP 3 CATIONS
Introduction
The group 3 cations that we will consider are Fe3+, Ni2+, Mn2+, and Al3+. These cations do not precipitate as insoluble
chlorides (as do those of group 1) or as sulfides in acidic solutions (like those of group 2). These ions can be separated from
those of group 4 by precipitation as insoluble hydroxides or sulfides under slightly alkaline conditions. The separation is shown
in the flow chart. As in group 2, the pH of the solution controls the sulfide-ion concentration. The slightly alkaline conditions
employed here favor a higher sulfide-ion concentration than that used in the group 2 separation. FeS, NiS, and MnS are more
soluble than the sulfides of group 2 and therefore require a higher sulfide-ion concentration for their precipitation. Making the
solution slightly alkaline reduces the hydrogen ion concentration. We see that decreasing the hydrogen ion concentration
causes the following equilibrium to shift to the right:
H2S(aq) === 2 H+(aq) + S2-(aq)
This results in an increase in the sulfide ion concentration.
Aqueous ammonia is a weak base:
NH3(aq) + H2O(l) === NH4+(aq) + OH-(aq)
A mixture of aqueous NH3 and NH4Cl makes a buffer solution whose pH allows the precipitation of FeS, MnS, and NiS.
Moreover, in this slightly alkaline solution, the insoluble hydroxides Fe(OH)3 and Al(OH)3 also to precipitate. Although Mg(OH)2 is
insoluble, it does not precipitate with the group 3 hydroxides because the OH- ion concentration is too small. The common ion
NH4+, from NH4Cl controls the OH- concentration and keeps it sufficiently low preventing the Mg(OH)2 from precipitating.
As you analyze the group 3 ions, pay particular attention to the colors of the solutions and precipitates. Aqueous
solutions of Al3+ ions are colorless while those Fe3+ appear yellow to reddish-brown; Ni2+ solutions are green, and Mn2+, very faint
pink.
The reactions involved in the precipitation of the group 3 cations are:
Fe3+(aq)
Al3+(aq)
3+
Ni(NH3)6 (aq) +
Mn2+(aq)
+ 3 OH-(aq)  Fe(OH)3 (s)
+ 3 OH-(aq)  Al(OH)3 (s)
S2-(aq)  6NH3 (aq) + NiS(s)
+ S2-(aq)  MnS (s)
red-brown
colorless
black
salmon
[1]
[2]
[3]
[4]
In aqueous NH3 solution, Ni2+ ions exist as the ammine complex Ni(NH3)62+, and addition of (NH4)2S results in the precipitation of
NiS. Some Fe(OH)3 is reduced by sulfide:
2Fe(OH)3 (s) + 3S2-(aq)  6OH- (aq) + 2FeS(s)
black
[5]
The group 3 precipitates are dissolved using HCl. Nitric acid is also added to help dissolve NiS by oxidizing S2- ions to
elemental sulfur. At the same time, nitric acid oxidizes Fe2+ to Fe3+:
Fe(OH)3(s) + 3H+(aq)  Fe3+(aq) + 3H2O (l)
FeS(s) + 2H+(aq)  Fe2+(aq) + H2S (g)
3Fe2+(aq) + NO3_(aq) + 4H+(aq)  3Fe3+(aq) + NO(g) + 2H2O (l)
Al(OH)3(s) + 3H+(aq)  Al3+(aq) + 3H2O (l)
3NiS(s) + 2NO3_(aq) + 8H+(aq)  3Ni2+(aq) + 3S(s) + 2NO(g) + 4H2O (l)
MnS(s) + 2H+(aq)  Mn2+(aq) + H2S (g)
[6]
[7]
[8]
[9]
[10]
[11]
After all the ions are in solution, addition of excess strong base allows the separation of Fe3+, Ni2+, and Mn2+ ions from Al3+ ions.
Fe(OH)3, Ni(OH)2, and Mn(OH)2 precipitate, while Al(OH), being amphoteric, redissolves in excess base forming aluminate ions:
CHEMISTRY OF GROUP 3 CATIONS
Fe3+(aq)
Ni2+(aq)
Mn2+(aq)
Al3+(aq)
+ 3OH-(aq)  Fe(OH)3 (s)
+ 2OH-(aq)  Ni(OH)2 (s)
+ 2OH-(aq)  Mn(OH)2 (s)
+ 3 OH-(aq)  Al(OH)3 (s)
- 27 -
red-brown
green
tan
colorless
[12]
[13]
[14]
[15]
The hydroxide precipitates are dissolved by the addition of H2SO4 to give a solution of Fe3+, Ni2+, and Mn2+ ions. After
the solution is divided into three equal parts, tests for the individual ions are made as described below.
Iron. A very sensitive test for Fe3+ uses the thiocyanate ion, SCN-. If Fe3+ is present, a blood-red solution results when SCN- is
added:
Fe3+(aq) + 3 SCN-(aq)  Fe(SCN)63-(aq)
red
[16]
Manganese. Because of its purple color, the permanganate ion, MnO4-, affords a suitable confirmatory test for Mn2+. When a
solution of Mn2+ is acidified with HNO3 and then treated with sodium bismuthate, NaBiO3, Mn2+ is oxidized to MnO4-:
2Mn2+(aq) + 5NaBiO3(s) + 14H+(aq)  5Bi3+(aq) + 5Na+(aq) + 7H2O(l) + 2MnO4-(aq)
purple
[17]
Nickel. The presence of Ni2+ is confirmed by the formation of a bright red precipitate when an organic compound called
dimethylglyoxime (abbreviated as H2DMG) is added to an ammoniacal solution:
Ni(NH3)62+(aq) + H2DMG (aq)
 4NH3(aq) + 2NH4+ (aq) + Ni(HDMG)2(s)
red
[18]
Aluminum. When a solution that contains Al(OH)4- is acidified and then made slightly alkaline with a weak base NH3, Al(OH)3
precipitates:
Al(OH)4- + 4H+(aq)  Al3+(aq) + 4H2O (l)
Al3+(aq) + 3NH3(aq) +
3H2O(l)  3NH4+ (aq)
+ Al(OH)3(s)
[19]
[20]
Aluminum hydroxide is not easily seen, for it is gelatinous, translucent substance. To help see the hydroxide it is precipitated in
the presence of a red dye. The dye is adsorbed on the Al(OH)3, giving it a cherry-red color.
Procedure:
G3-1. Precipitation of Group 3 Cations
Place the decantate from Procedure G2-1 or 7 drops of known or unknown solution in a small (10mm x 75mm) test tube. If the
decantate from Procedure G2-1 has a precipitate, centrifuge, decant, and discard the precipitate. Add 5 drops of 2M NH4Cl and
stir; add 15M NH3 dropwise with stirring, until the solution is just basic. Usually this requires only a few only a few drops of the
NH3. Then add 2 additional drops of the 15M NH3 and 1 mL (20 drops) of water. Stir thoroughly. Next add 10 drops of (NH4)2S
and mix thoroughly. Heat the test tube in a boiling-water bath for about 5 min. If excessive frothing occurs, temporarily remove
the test tube from the hot-water bath. Centrifuge and test for completeness of precipitation using 1 drop of (NH4)2S. Note the
color of the precipitate. Decant and save the decantate for group 4analysis, or discard, as appropriate.
Wash the precipitate two times with 20 drops of a solution made by mixing equal portions of water and 1M NH4C2H3O2. For each
washing, stir the precipitate with the wash solution and heat the mixture in the water bath before centrifuging. Discard the
supernatant wash liquid.
G3-2. Dissolution of Group 3 Precipitate
Treat the precipitate from Procedure G3-1 with 12 drops of 12M HCl, and cautiously add 5 drops of 16M HNO3 and carefully mix
the solution (CAUTION: HNO3 can cause severe burns. If you come in contact with the acid, wash the area with copious
amounts of water). Heat the test tube in a hot-water bath until the precipitate dissolves and a clear, but not necessarily colorless,
solution is obtained. Add 10 drops of water, centrifuge to remove any sulfur that has precipitated, and decant into a casserole.
Note the color of the decantate.
CHEMISTRY OF GROUP 3 CATIONS
- 28 -
G3-3. Separation of Iron, Nickel, and Manganese from Aluminum
Make the solution in the casserole from Procedure G3-2 strongly basic, using 6M NaOH, mixing thoroughly. If the precipitate is
pasty and nonfluid, add 12 drops of water. Note the color of the precipitate. Transfer into a test tube and centrifuge. Decant,
saving the decantate, which may contain aluminum, for Procedure G3-7. To the precipitate, add 20 drops of water and 10 drops
of 6M H2SO4. Stir and heat in a water bath for 3 minutes or until the precipitate dissolves (it may take a long time to dissolve).
Add 12 drops of water and divide the solution into 3 approximately equal volumes.
G3-4. Test for Fe3+.
To one of the three samples from Procedure G3-3, add 2 drops of 0.2M KSCN (potassium thiocyanate). A blood-red solution
confirms the presence of iron as Fe(SCN)63-. Traces of iron that have been introduced as impurities along the way will give a
weak test. If you are in doubt about your result, perform this test on 10 drops of your original sample.
G3-5. Test for Mn2+
To the second portion of a solution from Procedure G3-3, add an equal volume of water and 4 drops of 3M HNO3. Mix, and then
add a few grains of solid sodium bismuthate, NaBiO3. Mix thoroughly with a stirring rod and centrifuge. A pink or purple color is
due to MnO4- and confirms the presence of manganese.
G3-6. Test for Ni2+
To a third portion of the sample from Procedure G3-3 add 6M NH3 until the solution is basic. If a precipitate forms, remove it by
centrifuging and decanting, keeping the decantate. Add about 4 drops of dimethylglyoxime reagent, mix, and allow to stand.
The formation of a strawberry-red precipitate indicates the presence of nickel.
G3-7. Test for Al3+
Treat only half of the decantate from Procedure G3-3 with 16M HNO3 until the solution is slightly acidic. (CAUTION: HNO3 can
cause severe burns. Wash with water immediately if you come in contact with the acid.) Then add 15M NH3 while stirring until
the solution is distinctly alkaline. Allow at least 1 min for the formation of Al(OH)3. Centrifuge and carefully remove the
supernatant liquid with a capillary pipet without disturbing the gelatinous Al(OH)3. Discard the supernate. Wash the precipitate
two times with 20 drops of hot water, discarding the decantate. Dissolve the precipitate in 7 drops of 3M HNO3. Add 3 drops of
aluminon reagent, which colors the solution; stir; and add 6M NH3 dropwise until the solution is just alkaline (avoid an excess).
Stir and centrifuge. The formation of a cherry-red precipitate, not solution, confirms the presence of aluminum.
Review Questions
1. What are the symbols and charges of group 3 cations?
2. What are the colors of the following ions in aqueous solution: Fe3+, Al3+¸Ni2+?
3. What are the colors of the following solid: Fe(OH)3, MnS, Al(OH)3, Ni(OH)2?
4. How can Fe3+ be separated from Al3+?
5. How can Ni(OH)2 be separated from Al(OH)3?
6. Complete and balance the following equations:
a. Fe3+(aq) + 3 OH (aq) 
b. Al(OH)3(s) + 3H+(aq) 
c. FeS(s) + 2H+(aq) 
d. 3NiS(s) + 2NO3 (aq) + 8H+(aq) 
7. Give the formula for a reagent that precipitates
a. Pb2+ but not Ni2+
b. Fe3+ but not Al3+
8. What cation forms a
a. Blood-red solution with thiocyanate ion?
b. Bright red precipitate with dimethylglyoxime?
9. If solid NH4Cl is added to 3M NH3, does the pH increase, decrease or remain the same? Explain your answer.
CHEMISTRY OF GROUP 3 CATIONS
- 29 -
Group 3 Flow scheme.
7
CHEMISTRY OF GROUP 3 CATIONS
Data Sheet
Name: _____________________________
Date: ______________________
Part I: Known
Procedure
Reagent/s
Observations
Equations
Teacher’s
Sig.
CHEMISTRY OF GROUP 3 CATIONS
- 31 -
Part II: Unknown
Procedure
Reagent/s
Observations
Equations
Ions Present: ___________________________________________________
Ions Absent: __________________________________________________________
+/-
Experiment 8
CHEMISTRY OF GROUP 4 CATIONS
Introduction
In addition to the ammonium ion, the cations of group 4 consist of ions of the alkali and alkaline earth metals. The
cations we will consider in this group are Ba2+, Ca2+, NH4+¸ and Na+. Because their chlorides and sulfides are soluble, these ions
do not precipitate with groups 1, 2, or 3.
Sodium ions are a common impurity and were even introduced (as was ammonium ion) in some of the reagents that
were used in the analysis of groups 1, 2, and 3. Hence, in the analysis of a general unknown mixture, tests for these ions must
be made on the original sample even before performing the group analysis.
Barium. Because barium chromate, BaCrO4 (Ksp = 1.2 x 10-10), is less soluble than calcium chromate, CaCrO4 (Ksp = 7.1 x 10Ba2+ can be separated from Ca2+by precipitation as an insoluble yellow chromate salt:
4),
Ba2+(aq) + CrO42-(aq)  BaCrO4(s)
yellow
[1]
BaCrO4 is insoluble in the weak acid HC2H3O2, but it is soluble in the presence of the strong acid HCl, because it is the salt of a
weak acid, H2CrO4. After BaCrO4 is dissolved in HCl, a flame test is performed on the resulting solution. A green-yellow flame is
indicative of Ba2+. Further confirmation of Ba2+ is precipitation of BaSO4, which is white.
Ba2+(aq) + SO42-(aq)  BaSO4(s)
white
[2]
Calcium. Calcium oxalate, CaC2O4, is very insoluble (Ksp = 4.0 x 10-9). The formation of a white precipitate when oxalate ion is
added to a slightly alkaline solution confirms the presence of Ca2+:
Ca2+(aq) + C2O42-(aq)  CaC2O4(s)
white
[3]
Additional evidence for the calcium ion is obtained from a flame test. Dissolution of CaC2O4 with HCl, followed by a flame test,
produces an orange-red flame character that is characteristic of calcium ions.
Sodium. Most sodium salts are soluble. The simplest test for sodium ion is a flame test. Sodium salts impart a characteristic
yellow color to a flame; the test is very sensitive, and because of the prevalence of sodium ions, much care must be exercised to
keep equipment clean and free from contamination by these ions.
Ammonium. The ammonium ion, NH4+, is the conjugate acid of the base ammonia, NH3. The test for NH4+ takes advantage of
the following equilibrium:
NH4+(aq) + OH-(aq) === NH3(aq) + H2O(l)
[4]
Thus, when a strong base is added to a solution of an ammonium salt and this solution is heated, NH3 gas is evolved. The NH3
can easily be detected by its effect upon red litmus.
Procedure:
The unknown is a solid. Divide the unknown solid into two equal parts; keep one half for the flame test. The other half should be
dissolved with 10 drops distilled water.
G4-1. Separation and Identification of Ba2+
If the solution, “known,” or “unknown” contains only cations of group 4, place 7 drops of the solution in a small test tube;
otherwise, use the supernate from group 3 analysis, Procedure G3-1. Add 8 drops of 6M drops of 6M acetic acid, HC2H3O2, and
1 drop of 1M K2CrO4 and mix. The formation of a yellow precipitate indicates the presence of Ba2+. Centrifuge, saving the
CHEMISTRY OF GROUP 4 CATIONS
- 33 -
decantate for Procedure G4-2 to test for calcium. Dissolve the precipitate with 6M HCl and perform a flame test as described
below.
To perform the flame test, obtain a piece of platinum or Nichrome wire that has been sealed in a piece of glass tubing.
Clean the wire by dipping it in 12M HCl that is contained in a small test tube and heat the wire in the hottest part of your Bunsen
burner flame. Repeat this operation until no color is seen when the wire is placed in the flame. Several cleanings will be
required before this is achieved. Then dip the wire into the solution to be tested and place the wire in the flame. A pale green
flame confirms the presence of Ba2+. If the concentration of Ba2+ is very low, you may not detect the green color.
As further confirmation of barium, add 10 drops of 6M H2SO4 to the solution on which the flame test was performed. A
white precipitate confirms the presence of Ba2+.
G4-2. Test for Ca2+
Make the decantate from Procedure G4-1 alkaline to litmus with 15M NH3. If a precipitate forms, centrifuge and discard the
precipitate. Add 7 drops of 1M K2C2O4 (potassium oxalate) and stir. The formation of a white precipitate indicates the presence
of calcium ions. Should no precipitate form immediately, warm the test tube briefly in the hot-water bath and then cool.
Additional evidence for Ca2+ is obtained from a flame test. Dissolve the precipitate in 6M HCl and then perform a flame
test. A transitory red-orange color that appears when the wire is first placed in the flame and later reappears somewhat more red
as the wire is heated is characteristic of the calcium ion. If the concentration of Ca2+ is very low, you may not observe the red
color.
G4-3. Test for Na+
The flame test for sodium is sensitive, and traces of sodium ion will impart a characteristic yellow color to the flame. Just about
every solution has a trace of sodium and thus will give a positive test. On the basis of the intensity and duration of the yellow
color, you can decide whether Na+is merely a contaminant or present in substantial quantity. Using a clean wire, perform a flame
test on your original (untreated) unknown. To help you make a decision to the presence of sodium, run a flame test on distilled
water and then on a 0.2M NaCl solution. Compare the tests.
G4-4. Test for NH4+
Place 2 mL of the original (untreated) unknown or known in a100-mL beaker (or casserole) and add 2 mL of 6M NaOH. Moisten
a piece of red litmus paper with water and stick it to the convex side of a small watch glass. Cover the beaker with the watch
glass convex side down. (The litmus paper must not come into contact with any NaOH.) Gently warm the beaker with a small
burner flame; do not boil. Allow the covered beaker to stand 3 min. A change in the color of the litmus paper from red to blue
confirms the presence of ammonium ion.
CHEMISTRY OF GROUP 4 CATIONS
Review Questions
1. What are the symbols and charges of group 4 cations?
2. Which is less soluble: BaCrO4 or CaCrO4? What chemical property do you use to decide?
3. What is the color of:
4.
5.
6.
a. BaSO4
b. BaCrO4
c. CaCrO4
What color do the following impart to the flame?
a. Ba2+
b. Ca2+
c. Na+
What reagent will precipitate
a. Cu2+ but not Ba2+?
b. Ag+ but not Ca2+?
c. Ba2+ but not NH4+?
Complete and balance the following equations:
a. Ba2+(aq) + CrO42-(aq)
b. NH4+(aq) + OH-(aq)
- 34 -
CHEMISTRY OF GROUP 4 CATIONS
- 35 -
Group 4 Flow scheme.
8
CHEMISTRY OF GROUP 4 CATIONS
Data Sheet
Name: _____________________________
Date: ______________________
Part I: Known
Procedure
Reagent/s
Observations
Equations
Teacher’s
Sig.
Observations
Equations
+/-
Part II: Unknown
Procedure
Reagent/s
Ions Present: ___________________________________________________
Ions Absent: __________________________________________________________
Experiment 9
CHEMISTRY OF ANIONS
Introduction
A systematic scheme based upon the kinds of principles involved in cation analysis can be designed for the analysis of
anions. This would involve separation of the anions followed by their identification. However, it is generally easier to take
another approach to the identification of anions. An effort is made either to eliminate or verify the presence of certain anions on
the basis of the color and solubility of the samples; then the material being analyzed is subjected to a series of preliminary tests.
From the results of the preliminary tests and observations, certain of the anions may be definitely be shown to be present or
absent; then specific tests are performed for those anions not definitely eliminated in the preliminary tests and observations. The
preliminary tests include treating the solid with concentrated sulfuric acid and using silver nitrate and BaCl2 as precipitating
agents. However, in this class, no preliminary tests will be performed. We will go directly into the specific tests.
In this experiment the following 10 anions are considered: sulfate (SO42-), nitrate (NO3-), carbonate (CO32-), chloride
(Cl ), bromide (Br-), iodide (l-), chromate (CrO42-), phosphate (PO43-), sulfide (S2-), and sulfite (SO32-).
As in the case of cation identification, the physical and chemical properties of the compounds formed by the anions
provide the basis for their identifications.
Procedure:
The color of a substance offers a clue to its constituents. For example, many transition metal salts are colored. Those
of nickel, Ni2+; are generally green; of iron (III), Fe3+, reddish-brown to yellow, of iron (II), Fe2+, grayish-green; of Cr (III), Cr3+,
green to bluish-gray to black; of copper (II), Cu2+, blue to green to black; of cobalt (II), Co2+, wine-red to blue; and of manganese
(II), Mn2+, pink to tan. By contrast, only a few anions are colored, and these contain transition metals as well. The colored anions
are chromate (CrO42-), yellow; dichromate (Cr2O7), orange-red; and permanganate (MnO42-), violet-purple. Hence, the color of
the solid may be used as a very good indicator of the presence of these ions. Observe the color of your unknown and note it on
the report sheet.
1. SO42- Place 10 drops of a solution of the anion unknown in a test tube, acidify with 6M HCl, and add a drop of BaCl2 solution.
A white precipitate of BaSO4 confirms SO42- ions (see Note 1).
Ba2+(aq)
+ SO42-(aq)  BaSO4(s)
Note 1: Sulfites are slowly oxidized to sulfates by atmospheric oxygen. Consequently, sulfites commonly show positive test for
sulfates.
SO32-(aq) + O2 (g)  SO42-(aq)
SO42-(aq) + Ba2+(aq) 
BaSO4(s)
2. SO32- Place 10 drops of a solution of the anion unknown in a test tube, acidify with 6M HCl, add 2-3 drops of 0.2 M BaCl2,
and mix thoroughly. If a precipitate (BaSO4) forms, remove it by centrifuging and decanting. To the clear decantate, add 1 drop
of 3-percent H2O2 (see Note 2). The formation of a white precipitate of BaSO4 confirms SO32- ions.
Note 2: H2O2 oxidizes sulfite to sulfate.
SO32-(aq)
SO42-(aq)
+ H2O2(aq)
+ Ba2+(aq)
 SO42- (aq) + H2O(l)

BaSO4(l)
Any SO42- originally present was previously removed as BaSO4 by centrifugation.
3. CrO42- Place 2 drops of a solution of the anion unknown in a test tube, add 10 drops of water, and make the solution just
acidic with 3M HNO3. Add 5-6 drops of ether (KEEP THE ETHER AWAY FROM FLAMES) and a drop of 3 percent H2O2, stir
CHEMISTRY OF ANIONS
- 38 -
well, and then allow the precipitate to settle. A blue coloration of the top ether layer (see Note 3) confirms CrO42- ions (see Note
4).
Note 3: The blue coloration in the ether layer is due to the presence of chromium peroxide, CrO5.
2 CrO42-(aq) + 2H+(aq)  Cr2O72-(aq) + H2O(l)
Cr2O72-(aq) + 4H2O2(l) + 2H+(aq)  CrO5(aq) +
5H2O(l)
Note 4: If your anion unknown is not colored or did not indicate CrO42- in the H2SO4 reaction, or if both
conditions apply, this test may be omitted.
4. I- Place 5 drops of a solution of the anion unknown in a test tube, add 5 drops of 6M HC2H3O2, and
then add 2 drops of 0.2 M KNO2. A reddish-brown coloration due to the presence of I2 confirms I-. If the
brown color is very faint, add a few drops of mineral oil, and shake well. A violet color in the top (mineral
oil), layer confirms I-.
NO2-(aq) + H+(aq)  HNO2(aq)
+ 2I- (aq) + 2H+(aq) 
2HNO2(aq)
2NO(g)
+ I2(g) + 2H2O(l)
5. Br- Iodides will interfere with this test (Note 5). Place 5 drops of a solution of the anion unknown in a
test tube and add 5 drops of chlorine water. A brown coloration due to the liberation of Br2 confirms Br-. If
the solution is shaken with a few drops of mineral oil, the brown color will concentrate in the top layer which
is mineral oil. Allow about 20s for the layers to separate.
Cl2(aq)
+
2Br-(aq)

2Cl-(aq)
+
Br2(aq)
Note 5: Iodide ions, if present, must be removed before testing for bromide. To remove iodide, acidify the
solution with 3M HNO3 and add 2M KNO2, dropwise, with constant stirring, until there is no further increase
in the depth of the brown color. Extract once by shaking with 5 drops of mineral oil. Discard the mineral oil
layer. Boil the water layer carefully until the iodine has bee largely driven off. Test the colorless, or nearly
colorless, solution for bromide as directed above.
6. NO3 - Iodides, bromides, and chromates interfere with the test and must be removed (see Note 6) if they
are present. Place 10 drops of a solution of the anion salt in a small test tube, add 5 drops of FeSO4
solution, and mix the solution. Carefully, without agitation, pour concentrated H2SO4 down the inside of the
test tube so as to form two layers. Allow to stand for 1 or 2 min. The formation of a brown ring between the
two layers confirms NO3- ions. (CAUTION: Do not get the H2SO4 on yourself or on your clothing. If you
do, wash immediately with copious amounts of water.)
(1) 3Fe2+(aq) + NO3-(aq) + 4H+(aq)  3Fe3+(aq) + NO(aq) + 2H2O(l)
(2) NO(aq) + Fe2+(aq) (excess)  Fe(NO)2+(aq)
(brown)
Note 6: Iodides and bromides react with concentrated H2SO4 to liberate I2 and Br2.
SO42-(aq)
SO42-(aq)
SO42-(aq)
+
+
+
8I-(aq) +
2I-(aq) +
2Br-(aq) +
10H+(aq)
4H+(aq)
4H+(aq)



H2S(g) + 4I2(aq) + 4H2O(l)
SO2(g) + I2(aq
+ 2H2O(l)
SO2(g) + Br2(aq) + 2H2O(l)
CHEMISTRY OF ANIONS
- 39 -
Chromate ions, if present, will be reduced by Fe2+ to green Cr3+
2CrO42-(aq) +
2H+(aq) 
2+
+ 6Fe (aq) + 14H+(aq)

Cr2O72-(g)
Cr2O72-(g) +
H2O(l)
3+
3+
2Cr (aq) + 6Fe (aq) + 7H2O(l)
The colors of I2, Br2, and Cr3+ will interfere with detection of the brown color of Fe(NO)2+. Consequently, I-, Br-, and CrO42- must
be removed as follows: Place 4 drops of the unknown anion solution in a test tube and add 0.2M Pb(C2H3O2)2 until precipitation
is complete. Centrifuge and decant, discarding the precipitate (PbCrO4). Treat the decantate with 0.1M Hg(C2H3O2)2 until
precipitation is complete. Centrifuge the solution to separate the precipitate (HgI2, HgBr2) and treat the decantate as above to
test for nitrate.
7. CO32- Sulfites will interfere with this test for carbonates.
A. When sulfites are absent. Place a small amount of the solid anion unknown in a small test tube and add a few
drops of 6M H2SO4. If a colorless, odorless gas is evolved, hold a drop of Ba(OH)2 solution over the mouth of the
test tube using either an eyedropper of a Nichrome wire loop; CO32- ions are confirmed if the drop turns milky.
(1) 2H+(aq)
(2) CO2(g)
+ CO32-(aq)  CO2(g) + H2O(l)
+ Ba(OH)2 (aq)  BaCO3(s) + H2O(l)
B. When sulfites are present. Place a small amount of the solid anion unknown in a test tube and add an equal
amount of solid Na2O2. (CAUTION! Sodium peroxide is a very strong oxidizing agent and all contact with skin
should be avoided. If you do get some on yourself, immediately wash it off with large volumes of water.) Add 3-4
drops of water, mix thoroughly, then proceed as above to test for carbonate.
8. S2- Place a small quantity of the solid unknown anion in a test tube and (UNDER THE HOOD) add 10 drops of 6M HCl. Hold
a piece of filter paper which has been moistened with 0.2M Pb(C2H3O2)2 over the mouth of the test tube so that any gas which
escapes comes into contact with the paper. A brownish or silvery black stain (PbS) on the paper confirms the presence of S2-. If
no blackening of the lead acetate occurs after 1 min, heat the tube gently; if still no reaction occurs, add a small amount of
granulated zinc to the contents of the tube. If the lead acetate is not darkened, S2- is absent.
S2-(aq) + 2H+(aq)
Pb2+(aq) + H2S(aq)
Zn(s)
+
HgS(aq)
+
 H2S(g)
 PbS(s) + 2H+(aq)
2H+(aq)  Zn2+(aq) + Hg(l)
+ H2S(l)
9. ClSulfides, bromides, and iodides will interfere with this test (see Note 7). Place 10 drops of a solution of the anion
unknown in a test tube and add a drop of AgNO3 solution. A white, curdy precipitate confirms Cl- ions.
Ag+(aq)
+
Cl-(aq)
 AgCl(s)
Note 7: Since Ag2S, AgBr, and AgI are also insoluble in acid solution, this test is not conclusive unless S2-, Br-, and I- are
definitely shown to be absent. Chromate ions, if present in high concentrations, may also interfere.
Interfere from chromate can be eliminated by dilution with 3M HNO3. Sulfide ions can be removed by boiling the
solution after adding 2 drops of 6M H2SO4 until the escaping vapors give no test for H2S with lead acetate paper. If you suspect
the presence of chloride, or bromide, or iodide, or any combination of these, they may all be confirmed as follows:
Place 10 drops of a solution of the anion unknown in a test tube and add 5 drops of AgNO3 solution. Centrifuge
and discard the supernate. To the precipitate add 10 drops of concentrated NH3 solution and 3 drops of yellow ammonium
sulfide solution. Stir the mixture with a glass rod and warm gently until the black Ag2S coagulates. Centrifuge and discard the
precipitate. Transfer the solution to a 50-mL beaker and boil to expel NH3 and to decompose the ammonium sulfide. When the
solution becomes cloudy, add 5 or 6 drops of 6M HNO3 and continue heating until H2S is completely removed. Then carry out
the following tests:
A.
Place 1 drop of the solution on a piece of filter paper; add 1 drop of iodide-free starch solution and 1 drop of 0.2M
KNO2. A blue color confirms iodide.
CHEMISTRY OF ANIONS
- 40 -
AgI(s)
2Ag(NH3)2+(aq)
I-(aq)
B.
+
+
+
HNO2 (aq)
2NH3(aq)  Ag(NH3)2+(aq) + I-(aq)
(NH4)2S (aq)  Ag2S(s) + 2NH4+ (aq) + 4NH3(g)
+ 2H+(aq)  2NO(g) + I2- (aq) + 2H2O(aq)
I2 + starch  starch  I2 complex (blue)
If iodide is present, add 3 or 4 drops of 0.2M KNO2 to the solution in the beaker and boil until no more brown
fumes are evolved. If iodide is absent, proceed directly with the test for bromide. Cool the beaker to room
temperature by running cold water over its outer surface. To the room-temperature beaker, add 4 or 5 drops of
3M HNO3, then a small pea-sized piece of solid Na2O2. A brown coloration due to the presence of Br2 confirms
Br-.
2Br-(aq)
+
O22-(aq)
+
4H+(aq)  Br2 (aq) + 2H2O(l)
C. If bromide is present, boil the uncovered contents of the beaker for 30 s to expel the remainder of the bromine.
Allow to stand for 30 s and decant the solution into a test tube. Centrifuge if necessary. Add 5 drops of 0.1M
AgNO3. A white precipitate of AgCl confirms chloride.
Ag+(aq)
+
Cl-(aq)
 AgCl(s)
10. PO43A. In the absence of iodides. Place 4-5 drops of a solution of the anion unknown in a test tube, add 2 drops of 6M
HNO3, 3-4 drops of (NH4)MoO4 (ammonium molybdate), mix thoroughly, and heat almost to boiling for 2 min.
Formation, sometimes very slow, of finely divided yellow precipitate confirms phosphate.
PO43-(aq)
+ 12MoO42-(aq)
+
24H+(aq)
+ 3NH4+(aq)  (NH4)3PO4  12MoO4 (s) + 12H2O(l)
B. In the presence of iodide The phosphate test gives a green solution when iodide is present. If iodide is known to be
present, acidify the solution with 6M HNO3, add 4 drops of 0.1M AgNO3, centrifuge to remove AgI, and treat the
decantate as in 10A above.
Review Questions
1.
2.
3.
4.
5.
6.
What are the names and formulas of the anions to be identified?
Identify each of the following anions from the information given:
a. Its barium salt is insoluble in water but its copper salt is soluble.
b. Its copper salt is insoluble in water but its sodium salt is soluble.
c. Its mercury (II) salt is soluble in water but its mercury(I) salt is insoluble.
A mixture of barium and silver salts is water-soluble. What anions may not be present in this mixture?
A white solid unknown is readily soluble in water, upon tretament of this solution with HCl a colorless, odorless gas is
evolved. This gas reacts with Ba(OH)2 to give a white precipitate. What anion is indicated?
How would treatment with concentrated H2SO4 allow you to distinguish between the following:
a. BaSO4 and Hg(NO3)2?
b. CaBr2 and Na3PO4?
c. ZnS and BaSO4?
d. HgI2 and K2CrO4?
Five different salts each imparted a yellow color to a flame and reacted with cold H2SO4 as follows. Identify each salt.
a.
b.
c.
d.
e.
7.
8.
Effervescence was observed, and the evolved gas was colorless and odorless and did not fume in moist air.
Effervescence was observed, and the evolved gas was pale reddish-brown, had a sharp odor, and fumed strongly in moist air.
No effervescence was observed, and the solid changed color from yellow to orange.
Effervescence was observed, and the evolved gas was colorless, had a sharp odor, and fumed in moist air.
Effervescence was observed, and the evolved colorless gas had a very sharp odor but did not fume in moist air and did not discolor a
piece of filter paper which had been moistened with a solution of lead acetate.
How would you test for NO3- in the presence of I-?
How
would
you
test
for
CO32-
in
the
presence
of
SO32-?
9
CHEMISTRY OF ANIONS
Data Sheet
Name: _____________________________
Date: ______________________
Part I: Known
Anion
Reagent/s
Observations
Equations
Teacher’s
Sig.
CHEMISTRY OF ANIONS
- 42 -
Part II: Unknown
Anion
Reagent/s
Observations
Equations
Ions Present: ___________________________________________________
Ions Absent: __________________________________________________________
+/-
Experiment 10
ANALYSIS OF IRON BY VISIBLE SPECTROSCOPY
Objectives
You will analyze an unknown for iron by converting all the iron to Fe3+ and then to the blue colloidal dispersion of ferric
ferrocyanide, Fe[Fe(CN)6]3. The absorbance will be read at 713 nm and compared with values from a calibration curve. Each
student will receive an individual solid unknown, but will use standards prepared with three others in the class.
Reagents
20 ppm Fe3+stock solution,
Materials (per group)
2 of 50 mL buret, buret clamp, ring stand, wash bottle, 400 mL beaker, stock bottle, 2 of 50 and 1 of 250 mL volumetric
flasks, 9 sample bottles, 1 cuvette, hot plate.
Procedure
A. Preparation of Calibration Standard Solutions:
Each team of four students should get 100 mL of 20-ppm stock Fe3+ solution. Using one buret, and your 50-mL
volumetric flasks, prepare the following dilutions: 1, 2, 4, 6, and 8 ppm, adding 1 mL of 1% potassium ferrocyanide, K4Fe(CN)6
·3H2O, just before diluting the mix to 50 mL. These solutions are stored in small plastic bottles.
B. Preparation of blank:
Add 4.0 mL of 1:1 H2SO4, 1 mL of 1 % potassium ferrocyanide and one or two drops of 0.1 N KMnO4 (to make a
slight pink color) to your 50-mL volumetric flask and dilute to volume. Only one blank per team is allowed.
C. Preparation of Individual Unknown:
1.
Dissolve the unknown in 100 mL water in a 400 mL beaker. Add 4 mL of 1:1 H2SO4. Warm over a hot plate (to
approximately 50OC) and oxidize with approximately 0.1-N KMnO4 solution (by slowly adding dropwise using capillary pipet /
medicine dropper) until the iron solution remains faintly pink. Cool this, transfer completely to and dilute to mark in the 250 mL
volumetric flask. This is your unknown stock solution and may be stored in the volumetric flask for a few hours.
2.
Deliver from the other buret exactly 5.00 mL of your unknown stock solution, add 1 mL of 1% potassium ferrocyanide
and dilute to 50 mL in a volumetric flask. Do three aliquots in this way and store in small plastic bottles (triple wash bottles with
each solution before final storage).
D. Analysis of Standards, Unknown, and Blank:
Read the absorbance of each group standard, your sample (three aliquots) and blank at 713 nm at least 15 minutes
after final mixing, but shake it just before the reading.
For each reading, triple-wash the cuvette with the solution you are going to measure it with.
ANALYSIS OF IRON BY VISIBLE SPECTROSCOPY
- 44 -
Draw your Beer’s Law plot and determine the concentration of your unknown in ppm. If you use Microsoft Excel, use the Chart
feature to create the plot. After creating the plot, click Chart…Add Trendline to automatically draw the best-fit line. In Add
Trendline dialog box, select the ”Type” tab and slect type “Linear”. Then select the “Options” tab, check the “set intercept = 0”;
“display equation on chart”, and “Display R-squared on chart”.
Using the equation derived from the data, calculate the ppm Fe of your unknown.
Review Questions
1.
2.
3.
What is Beer’s Law?
What is the purpose of the blank?
What is the purpose of triple washing?
10
ANALYSIS OF IRON BY VISIBLE
SPECTROSCOPY
Data Sheet
Name: _____________________________
Date: ______________________
Calibration
Standards
Blank
1 ppm
2 ppm
4 ppm
6 ppm
8 ppm
Absorbance
Concentration
Unknown
Trial 1
Trial 2
Trial 3
Average
Absorbance
Concentration
Linear Regression equation: concentration
= m * (absorbance) + b
=_____ * Abs + _______
Concentration of the Unknown Fe Using Linear Regression: _______________________