Contents
Page
Guidance and Regulations for working in CHEMY101
Laboratories.
Safety and Lab Techniques.
2
Common Laboratory Apparatus.
10
Significant Figures.
13
Standard Deviation.
16
Experiment 1
The Equation for a Reaction.
18
Experiment 2
Limiting Reactant.
20
Experiment 3
Experiment 4
Experiment 5
Experiment 6
Using Dilution Factors in an Acid-Base Titration and
Analysis of Vinegar.
Gravimetric Analysis of a Chloride Salt.
Determination of the Molecular Mass of Metal Carbonate
By Back Titration.
Spectrophotometric Analysis of a Commercial Aspirin
Tablet.
6
23
27
32
34
Experiment 7
Chemical Reactions.
37
Experiment 8
Determining the Molar Mass of a Gas.
45
Experiment 9
Separation of a Mixture by Paper Chromatography.
50
1
Guidance and Regulations for working in CHEMY101
Laboratories.
1. Personal Safety Regulations
1.1 Safety Spectacles
Wearing safety spectacles or eye protection is obligatory, safety google must be worn at all
time, in all laboratories. Failure to wear eye protection in a laboratory will result in the person
being asked to leave immediately.
1.2 Protective Clothing
Laboratory coat is required to work in all laboratories. Wearing clothing that gives full coverage
of the legs and hands is important. Failure to obtain a laboratory coat will result that the person
being warned for the first time and expelled if the same case is repeated.
1.3 Footwear
It highly advised to wear shoes, rather than sandals when working in laboratories. This will
gives adequate protection from spillages of any chemicals.
1.4 Food and Drinks
Avoid accidental ingestion of substances by NEVER eating, drinking, smoking, using mouth
pipettes etc in the laboratory.
1.5 Washing Hands
Always wash your hands on leaving the laboratory for food or visiting the toilet, and at the end
every experiment.
1.6 Safety equipments
Always know where to locate the nearest eyewash fountain, emergency shower, fire
extinguisher, fire blanket and first aid kit.
1.7 Handling Chemicals.
Avoid skin contact with the substance you need to handle by using suitable apparatus if
necessary. If there is skin contact with any substance, wash your skin immediately under running
water.
2
1.8 Unauthorized Procedures.
Do not start new or unfamiliar procedures until the precautions have been discussed with your
instructor, and always ask if you in doubt. Never work alone at any time in the laboratory.
2. Laboratory Report.
* The report will be corrected out of 10 marks. 5 marks are given to the student’s work and 5 for
the laboratory report.
* The report has to include the following:
( Name, I.D#, Section No., Course No., Title of Experiment, No. of Experiment, Aim (0.5
mark), Introduction (0.5 mark), Chemical and glassware (Optional), Procedure (Optional),
Results and Calculations (2 marks), Discussion (1 mark) and Conclusion (1 mark)).

Cheating or copying from another report is strictly forbidden. If the instructor finds
any case of copying, the two student should be warned and 2 marks has to be taken from
both reports. If the same cases repeated again, the reports will not be accepted and the
students will get 5 out of 10 to his/her experimental work.

The late report will not be accepted after two weeks from finishing the experiment. If the
report is submitted within two weeks, the student will get 5 out 10 in that experiment.
3. Prelaboratory Quizzes.
There will be a prelaboratory quiz every week before the practical work begins about
the proposed experiment. The quiz duration is 10 minutes only and it must be held during
the first 15 minutes of the practical session. The student will not be allowed to take the quiz
if he/she attends 15minutes pass the hour of laboratory time. The corrected quizzes can be
shown to students and kept with the instructor afterward.
4. Laboratory Tests.
There will be two laboratory tests during the semester that are taken within the same time of
the 2nd and 3rd hour exams of theoretical part of the course. The tests will cover the
experiments taken equally.
3
5. Students who are repeating the course.
Those students have the choice not to repeat the practical part of the course if they passed the
practical part before. Any student who satisfies these conditions and decides not to take the
practical again, the student has to inform the course instructor about his/her decision and the
details of the time and marks of the practical when was taken for the first time. Those
students who will not repeat the practical again will not take the laboratory tests also.
6. Attendance
Attending all the practical sessions is extremely important and any student fails to do so will
face the following:

According to UOB regulations, a warning letter will be received if two laboratory
sessions are missed with or without excuses.

According to UOB regulations, a WF grade will be obtained if three laboratory
sessions are missed with or without excuses.
Any student who comes to the laboratory after 15 minutes pass the hour will not be allowed
to attend and he/she will be considered as an absent.
Any student who attends an alternative session during the same week if the laboratory is
missed, it is the student’s responsibility to inform the instructor about that.
7. The first laboratory session.
The first session of the practical is held in the laboratory without working in an experiment.
The following topics will be discussed:

The safety regulations followed in the laboratory. (see section 1 above).

The common glassware and apparatus. Brief information about the use of each glassware.

Handling solids and types of balances.

Significant figures.

Standard deviation.
4
8. The first Prelab quiz.
The first quiz will be held in the week followed the first laboratory session. It will include the
following:

Glassware

Significant figures

Standard deviation

Prelab quiz in experiment 1.
9. Assessments.
The practical part of the course has 15% of the total grades. The practical marks are distributed
as follows:

8% for working in the laboratory and writing the laboratory reports.

4% for the prelaboratory quizzes.

3% for the laboratory tests.
******Good luck ******
5
Safety and Lab Techniques
Safety in the Laboratory
Safety First
Safety Codes
Safety is of paramount importance in any Laboratory. It is all too
easy to lose your sight or damage your limbs and clothing.
Laboratories are dangerous if the safety regulations are not
observed, but are quite safe for mature students who observe, at
all times, the safety regulations given by the instructor.
Toxic
Explosive
Harmful
Irritant
Highly
flammable
Oxidising
Corrosive
Environmental
Toxicity
Safety Rules
The Safety rules for your laboratory will be given to you by your
instructor. They will include :
1.
Use of protective clothing.
2.
Forbidden attempts of unauthorized experiments.
3.
Asking , if in doubt. It is common sense to check on a
point, not an admission of weakness. This must never
be forgotten, no matter what the situation is.
A Check List of Safety Equipment
Make certain that you know the location of the following items of
safety equipments by drawing a plan of your laboratory and
placing in the details given below and any other requirements
given by your instructor.
1.
2.
3.
Location and use of fire extinguisher.
Location and use of eye wash.
Location of first aid kit.
4.
Location of emergency exists.
5.
Location and use of fume cupboards.
6.
Location and use of fire blanket.
DANGER
Biological
hazard
Glass hazard
6
Check List of Laboratory Hazards
Watch out for these hazards in a laboratory.
1.
2.
3.
4.
5.
6.
7.
8.
Poor ventilation of fumes.
Corrosive chemicals.
Cutting yourself with glass apparatus.
Thinking that because a chemical is often used that it is safe. For example, sodium
hydroxide solution, a common laboratory reagent, is hazardous to skin and a disaster to
the eyes!
Carelessness. Chemistry, like flying, is inherently safe but unforgiving of carelessness.
Haste.
Ignorance.
Any others given by your instructor.
HANDLING SOLID CHEMICALS
General Guidelines
Each laboratory has its own system for the handling of solid chemicals, so the advice given here
must always be related back to what your instructor expects of you. However, certain steps are
essential. These are :
1.
2.
3.
Read the label – have I got the right chemical ?
Transferring the chemical from the stock bottle in the store of the instructor bench to
your bench. One method is to use the spatula to get the required amount of chemical
from the bottle on to a piece of paper folded into a ‘V’. Do this at the service bench. Do
not take the chemical bottle to your bench.
Transferring the chemical to the test tube, or beaker, or other kit.
Beaker present no problems but narrow openings do. This is where the ‘V’ folded paper
becomes useful. Not only it gives some rigidity to the paper, but with gentle tapping you
can guide the chemical into the narrowest glassware and without spilling any!
Figure 1
4.
5.
6.
Do not use much more of the chemical than the instructions require. Chemical chaos
results.
Do not put waste solids into the sinks. Bins should be provided.
Do not mix chemicals and this includes:
(a.) Use a dirty spatula to remove a chemical from its stock bottle; the resulted
contamination will be significant.
(b.) Always ensure when returning unused chemical, that it goes back into the stock
bottle from where it came, and no other !
7
Balances
You will meet two types of balances during your practical work. One is for when the amounts of
solid chemical to be used are not critical and the other is for when very accurate weights are
necessary.
* Top loading balance
Figure 2
This balance is used for appropriate weighing. It will give you a value within + 0.1 g. Your
instructor will explain its use when required.
The analytical balance
This balance is accurate, fast, easy to use and also, very expensive. It may give a reading to within
+ 0.005 g and therefore, is used when a high degree of accuracy is required.
Steps in using the analytical balance
Because there are a number of designs of balances in use, your instructor will explain how to use the
balance in your lab. List the instructions given by your instructor in handling the analytical balance
in your notebook. Suggested headings for these notes might be :
Preparing balance for use
Weighing
Tarring
Closing down
Figure 3
8
Some don’ts in weighing
(1)
(2)
(3)
Do not place chemicals directly onto the pan. Always use an appropriate container for
chemicals. This is to prevent corrosion of the pan.
Do not place hot objects in the balance. They can blister the protective coating of the
pan and also set up convection currents, which move the pan and therefore, make
weighing difficult.
Do not move electric balance. For accurate readings they must be leveled. Moving a
balance will upset its level. If it is necessary to move a balance from one place to
another, it must be closed down while being shifted and then relevelled before being
used. Your instructor will demonstrate this if shifting is required.
9
Common Laboratory Apparatus
Beaker
Test tubes
Watch glass
Dropping bottle
Funnel
Florence (round bottom)
flask
Conical flask
10
Test tube block
Test tube holder
Clamp holder
Clamp
Pipette bulb
Tong
Burner
11
Spatula
Pipettes
Distilled water bottle
Buchner flask
Burette
Graduated cylinder
Volumetric flasks
Dropper
Buchner funnels
12
Stirring rods
Crucible and cover
SIGNIFICANT FIGURES
Numbers (or digits) that arise a result of measurement are called significant figures, which are
the meaningful digits in a measured quantity. It is easy to determine how many significant are
present in a number by using the following rules:
1. All non zero digits are significant.
e.g.: 2.1 has two sig. figs.
328 has three sig. figs.
Q. How many sig. figs, are in these numbers?
1) 29 _______
2) 3.14 _________
3) 32987 ___________
2. Zeros at the left of a measurement are not significant.
These zeros are used to indicate the placement of the decimal point.
e.g.: 0.0097 has only two sig. figs.
0.0195 has three sig. figs.
Q. How many sig. figs, are in these numbers?
1) 0.0001 _______
2) 0.0987_________
3) 0.0076 ___________
3. Zeros between non-zero digits are significant.
e.g.: 903 has three sig. figs.
1.04 has three sig. figs.
0.0204 has three sig. figs.
Q. How many sig. figs, are in these numbers?
1) 406 _______
2) 28.07_________
3) 1.002___________
4. Zeros at the right of a number, after a decimal point, are significant.
e.g.: 0.400 has three sig. figs.
500.0 has four sig. figs.
Q. How many sig. figs, are in these numbers?
1) 0.500 _______
2) 0.0800_________
3) 6.090___________
13
5. For numbers that do not contain decimal points, the zeros to the right of a non-zero digit
may or may not be significant.
e.g.: 400 ( the zero may be not significant) it can be written as:
4.00 × 102 three sig. figs.
4.0 × 102 two sig. figs.
4 × 102
one sig. figs.
Q. How many sig. figs, are in these numbers?
1) 3.600 × 102 _______
2) 5 × 10-6 _________
3) 6.000 × 104 ___________
6. In multiplying and dividing, an answer cannot have more sig. figs. than the measurements
with the least number of sig. figs. used.
e.g.: 48.22 × 1.64 = 79.1 (The answer must has 3 sig. figs., measurement with least sig. figs.)
52.7
 88 (The answer must has 2 sig. figs., the least precise measurement has 2 sig. figs.)
0.60
Q. Write the answer to the following problems with the correct number of sig. figs.
a) 5.237 × 6.4 = _________________
b)
0.5178
 ___________
0.087
7. The concept of precision is used only with the operations of addition and subtraction. The
answer cannot be more precise that the least precise measurements used to obtain the
answer.
e.g.:
28.14
+
10.2
Least precise measurement, answer must be rounded to the tenths place, i.e. it should
38.34
be given as 38.3 not 38.34
Q. Write the answer to the following problems with the correct number of sig. figs.
a)
b)
104.31
58.010
 22.111
 12.00
14
8. Numbers that are not measured are called exact numbers. Exact numbers do not have sig.
figs. and ignored when determining the number of sig. figs. in an answer.
e.g. 44.88/2 = 22.44 ( the answer has must four sig. figs. not one.)
9. Rounding off numbers.
A number is rounded off to a given place, when the digit immediately to the right of that place
is greater than 5. We add 1 to the place we are rounding to. If the digit is less than 5 we simply
eliminate that digit.
e.g.: If 2.489 is to be rounded to the tenth place, 2.5 is the correct answer. 0.0932 is to be
rounded to the hundredths place, 0.09 is the correct answer.
If precision is required in rounding off, and the digit immediately to the right is 5 and the number
to be rounded off is even, the 5 is eliminated, it the number is odd then 1 is added.
e.g.: 2.35 is rounded to 2.4
2.65 is rounded to 2.6
Q. Round off these numbers to the indicated place.
a). 6.429 to the tenth places. _______________.
b). 5.009 to the hundredth places_________________.
15
Standard Deviation
As a mean of estimating the precision of your results, it is desirable to calculate the
standard deviation. Before we illustrate how to do this, however, we will define some
terms.
Accuracy: measure how closely individual measurements agree with the
correct (true) value.
Precision: the closeness of agreement among several measurements of the
same quantity; the reproducibility of the measurement.
Error: difference between the true result and determined result.
Mean: arithmetic mean or average (μ), where

sum of results
number of results
The scatter about the mean–that is, the deviations from the mean–are measures of
precision. Thus the smaller the deviations, the more reproducible or precise the
measurements.
Standard deviation (s) is related to statistics and is a better measure of precision and is
calculated using the formula
s
sum of the squares of the deviations from the mean
number of observatio ns - 1
 i i  

N 1
2
Where s = standard deviation from the mean, χi = members of the set, μ = mean, and N =
number of members in the set of data.
Example:
An experiment's results are 1, 3, and 5. Calculate the mean, the deviations from the mean,
the standard deviation, and the relative standard deviation for the data.
The mean μ = 3.
The deviations from the mean are χi – μ= 2, 0, and 2.
Standard deviation s 
22  02  22
2
3 1
The results of the experiment could be reported as 3  2.
The relative standard deviation is 2/3 = 0.7, or 70%
16
Exercise 1:
The results of an experiment are 2.100, 2.110, and 2.105.
(a) Calculate the mean, the deviations from the mean, the standard deviation, and the
relative standard deviation.
(b) Write the results of the experiment in terms of the standard deviation.
(c) Are the data in Exercise 1 more or less precise than those in the example above?
and why?
(d) Are the more precise data necessarily more accurate?
The standard deviation may be used to determine whether a result should be retained or
discarded. As rule of thumb, you may discard any result that is more than two standard
deviations from the mean. For example, if you had result of 49.65% and you had
determined that your percentage of a given element was 49.25  0.09%, this result
(49.65%) could be discarded. This is because s = 0.09 and 49.25 – 49.65 = 0.40, which
is greater than 20.09. This result is more than two standard deviations from the mean.
Exercise 2:
The results of an experiment are 32.52%, 32.14%, 32.61%, and 32.72%.
(a) Find the mean, the standard deviation, and the relative standard deviation.
(b) Can any result be discarded?
17
EXPERIMENT 1
THE EQUATION FOR A REACTION
Aim:
a.
To determine the molar ratio between sodium hydrogen carbonate and hydrochloric
acid when they react.
To write an equation for the reaction.
b.
Introduction:
A chemical change is called a chemical reaction.
Chemical reactions are
represented by chemical equations. A chemical equation uses chemical symbols to show
what happens during a chemical reaction and to identify reactants and products.
A
chemical equation represents the overall change that takes place in a chemical reaction by
showing the number and kind of atoms, molecules or ions present before and after a
reaction. Formulas of reactants appear on the left side of the equation and those of products
are written on the right. In a balanced chemical equation, the number of atoms of a given
element is the same on both sides of the equation. The coefficients of a balanced equation
represent the relative number of moles of reactants and products. These coefficients should
be maintained as the simplest whole number ratio of reactants and products. A balanced
chemical equation illustrates the following:
*
The relative number of atoms, molecules and ions taking part in a chemical
reaction.
*
The relative number of moles of atoms, molecules and ions involved.
*
The mass relation between reactants and products.
*
The state of reactants and products by using:
(g) for gas
(l) for liquid
(s) for solid
(aq) for aqueous solution.
18
Procedure:
Preliminary Test (Gas test)
1. Set up the apparatus as shown.
Delivery tube
Lime water
HCl and NaHCO3
2. Half –full the test tube with lime water, Ca(OH)2, and place about 2 g of NaHCO3(s)
and 3 cm3 of 2 M HCl(aq) in the conical flask and stopper. Observe the reaction
between the evolved gas and the lime water. Record your observations and write a
chemical equation.
3. Repeat step 2 and collect the evolved gas in an empty test tube, then test with a
lighted splint.
Equation of a Reaction ( Mole ratio and salt test)
1. Weigh a dry evaporating dish and record the mass.
2. Add about 0.6 g of sodium bicarbonate, reweigh and record the mass.
3. Place the evaporating dish on a hot plate and carefully (dropwise) add dilute
hydrochloric acid (2 M). When most of the bicarbonate has reacted, begin to heat
the dish.
4. While heating, check that all the bicarbonate has reacted by adding more of the acid
drop by drop until no more effervescence is observed.
5. Heat to dryness, not heating too strongly because you may lose some of the solid
by spitting as the solid becomes dry.
6. Let the dish to cool to room temperature, weigh and record the mass.
7. Dissolve some of the residue in a test tube using distilled water. Add few drops of
silver nitrate to test for Cl- ions. Record your observations.
Calculation:
* Calculate the moles of NaHCO3 and moles of NaCl and find the ratio. Write a balanced
chemical equation for the reaction.
* Write the equation for the reaction between silver nitrate and Cl-.
Why is it a good way to test for Cl-?
19
EXPERIMENT 2
LIMITING REACTANT
Aim:
To determine the limiting reagent and percentage yield of a chemical reaction.
Chemical and other requirements:
0.125 M Sliver Nitrate ( AgNO3 ),
0.1 M Potassium Chromate ( K2CrO4 )
, Corrosive
Dangerous for Environment
25ml Graduated Cylinder, 250 ml Conical Flask, 250 ml Beaker, Glass Funnel,
Watch Glass, Droppers
Introduction:
When a chemist carries out a reaction, the reactants are usually not present in exact
stoichiometric amounts, that is, in the proportions indicated by the balanced equation. The
reactant used up first in a reaction is called the limiting reagent, since the maximum amount
of product formed depends on how much of this reactant was originally present (Figure 1).
When this reactant is used up, no more products can be formed. Thus one or more of the
other reactants will often be present in quantities greater than those needed to react with the
quantity of the limiting reagent present. These reactants are called excess reagents.
The concept of the limiting reagent is analogous to the relationship between the
number of stamps available and the number of letters to be mailed. If there are nine letters
and only six tamps, then the maximum number of letters that can be sent is six. The number
of stamps thus limits the number of letters that can be mailed, and there is an excess of
envelopes.
Figure 1.
Diagram showing the complete use of a limiting reagent (open circles) in a
reaction.
20
Often, you will be told the amounts of two different reactants and asked to determine
which is the limiting reactant and calculate the theoretical yield of product. To do this, it
helps to follow a systematic procedure. The one we will use involves three steps :
1.
2.
3.
Calculate the amount of product that would be formed if the first reactant was
completely consumed.
Repeat this calculation for the second reactant; that is, calculate how much product
would be formed if all of that reactant was consumed.
Choose the smaller of the two amounts calculated in (1) and (2). This is the theoretical
yield of product; the reactant that produces the smaller amount is the limiting reactant.
The other reactant is in excess; only part of it is consumed.
Procedure:
1.
2.
3.
4.
5.
Using graduated cylinders measure 20.0 ml of 0.125 M AgNO3 solution and 13.0 ml
of K2CrO4 solution.
Mix the two solutions together in a 250 ml beaker.
Leave the mixture to settle down. Meanwhile, weigh a filter paper labeled with your
initials.
Filter the solid on the labeled filter paper, put it on a watch glass and place it in the
oven to dry.
Weigh the solid and filter paper then re-dry in the oven and reweigh until you get a
constant weight.
Calculation:
1.
Determine the limiting reagent. The equation for the reaction is
2 AgNO3(aq) + K2CrO4(aq)  2KNO3(aq) + Ag2CrO4(s)
2.
3.
Calculate the theoretical yield.
Calculate the % yield.
Point for discussion:
1. Suggest reasons why the actual yield of a reaction is almost always less than the
theoretical yield.
21
Safety points for chemical used
Silver Nitrate ( AgNO3)
silver nitrate is an oxidant and should be properly stored away from organic compounds. silver
nitrate is toxic and corrosive. Brief exposure to the chemical will not produce immediate or even
any side effects other than the purple, brown or black skin stains, but with more exposure, side
effects will become more noticeable, including burns. Long-term exposure may cause eye
damage. Short contact can lead to deposition of black silver stains on the skin. Besides being
very destructive of mucous membranes, it is skin and eye irritant.
Potassium Chromate (K2CrO4)
Potassium Chromate is very toxic and may be fatal if swallowed. It may also act as a
carcinogen, and can create reproductive defects if inhaled or swallowed. It also is a strong
oxidizing agent if in the presence of H+ to produce the dichromate ion. It may react rapidly, or
violently. It is also possible that it may react explosively with other reducing agents and
flammable.objects.
22
EXPERIMENT 3
USING DILUTION FACTORS IN AN ACID-BASE TITRATION AND
ANALYSIS OF VINEGAR
Aim:
1. To prepare volumetric solutions by dilution of more concentrated solutions.
2. To determine an unknown concentration of sodium hydroxide solution
3. To acquire experience in applying solution stoichiometry rules to an acid – base
titration.
4. To determine the concentration of acetic acid in commercial vinegar.
Introduction:
This experiment begins with the dilution of more concentrated solutions of 1.00 M
hydrochloric acid and an unknown concentration of sodium hydroxide. The HCl solution of
known concentration will then be used to determine the concentration of the base, NaOH,
using phenolphthalein indicator.
A titration measures the volume of solution delivered from a buret. In this experiment
sodium hydroxide is titrated into a flask containing an acid. After a sufficient amount of base
is added to neutralize the acid in the flask, we will stop the titration. This is termed the
endpoint and is signaled by an indicator that changes color. The indicator used in this
experiment is phenolphthalein. Phenolphthalein is colorless in acid and red-pink in base. A
single drop of base is sufficient to bring about the color change.
Vinegar contains acetic acid, CH3CO2H, which is an organic acid with one acidic
hydrogen. The percent by weight of acetic acid in vinegar can be determined by titrating a
measured volume with standardized NaOH.
NaOH(aq) + CH3CO2H (aq) 
23
CH3CO2 Na(aq) + H2O(l)
Procedure:
(a.)
DILUTION:
Pipette 4.0 mL of the stock 1.00 M HCl solution into a 100.0 mL volumetric
flask. Add sufficient distilled water to make 100.0 mL of solution.
Calculate the morality of your solution.
The diluted HCl solution is
M
Pipette 10.0 mL of the concentrated sodium hydroxide solution into a 100.0
ml volumetric flask, add sufficient distilled water to make 100.0 ml of solution.
Dilution factor is ____________________
(b.)
Titration of HCl and NaOH
1. Clean a burette by rinsing it twice with small quantities of distilled water.
Empty the water by opening the stopcock of the burette. Then rinse the
burette with about 5 mL of the diluted sodium hydroxide solution. It is
always a good practice to make the last rinsing of any glassware with the
solution it will contain.
2. Fill the burette with the diluted NaOH solution. Be certain to open the
stopcock and allow a few mL of the solution to flow out. Check to see that
there aren’t any air bubbles between the tip of the burette and the stopcock.
3. Record the volume of the NaOH solution by reading the bottom of the
solution’s meniscus. Be certain to record this initial burette reading in your
data sheet.
4. Carefully, measure exactly 10.0 mL of the diluted HCl in 250 mL conical
flask using a pipette.
5. Add 2 or 3 drops of phenolphthalein indicator to the HCl solution and swirl
it thoroughly.
6. Titrate the HCl acid solution with the sodium hydroxide solution by opening
the stopcock of the burette in very slow rate. Be certain to swirl the HCl acid
solution during the titration. Continue adding the NaOH solution until the
HCl acid solution turns to a permanent pale pink color. At this point, stop
adding NaOH solution.
7. Record your final buret reading in your data sheet.
8. Repeat steps (4 to7) twice to have three volume readings of NaOH solution.
24
(c.)
Analysis of Vinegar
1. Pipette 5.0 mL of the stock of vinegar solution into a 100.0 mL volumetric
flask. Add sufficient distilled water to make 100.0 mL of solution and mix
the solution thoroughly.
2. Pipette 10.0 mL of the diluted vinegar solution to a 250-mL conical flask.
3. Add 2-3 drops of phenolphthalein indicator solution to the vinegar solution.
4. Titrate the vinegar solution with the same NaOH solution which is used in
part b (same burette).
5. Stop the titration when the vinegar solution turns to pale pink.
6.
Record your final burette reading in your data sheet.
7. Repeat steps (2 to 6) twice to have three volume readings of NaOH solution.
Calculation
1.
2.
Calculate the average volume of NaOH solution added .
The equation for the neutralization is :
HCl(aq) + NaOH(aq)
NaCl(aq) + H2O(l)
Calculate the number of moles of NaOH in your average volume.
3.
Calculate the molarity of the diluted NaOH solution. (Note : molarity = the number of
moles per litre of solution).
4.
From your dilution factor, calculate the molarity of the original NaOH stock solution.
5.
Calculate the number of grams/liters (g/L) of NaOH in the stock solution.
6.
Calculate the molarity of acetic acid in the diluted vinegar solution.
7.
Calculate the molarity of acetic acid in the vinegar stock solution.
8.
Calculate the concentration of the vinegar stock solution in (g/L) and mass percent.
Questions:
1.
2.
How many mL of water should be added to 25.0 mL of 6 M nitric acid to make 2M
solution?
What is the volume of 0.50 M H2SO4 required to neutralize 10.0 mL of 0.75 KOH
solution?
25
Table for Part b.
HCl
Titrant
(known
Volume of titrant
concentration)
NaOH
Analyte (Unknown
Indicator
concentration)
Titrations
Readings of the burette
Initial (I) mL
Final (F) mL
1
10.0 mL
Phenolphthalein
Volume used = F-I
(mL)
2
3
4
Average volume
Table for Part c.
Titrant
Analyte
Titrations
NaOH
Volume of analyte
CH3COOH
Indicator
Readings of the burette
Initial (I) mL
Final (F) mL
1
2
3
4
Average volume
26
10.0 mL
Phenolphthalein
Volume used = F-I
(mL)
Experiment 4
Gravimetric Analysis of a Chloride Salt
Objective:
To illustrate typical techniques used in gravimetric analysis by quantitatively
determining the amount of chloride in an unknown.
Introduction:
Quantitative analysis is that aspect of analytical chemistry concerned with determining
how much of one or more constituents is present in a particular sample of material.
Information such as percentage composition is essential to establishing formulas for
compounds. Two common quantitative methods used in analytical chemistry are
gravimetric and volumetric analysis. Gravimetric analysis derives its name from the fact
that the constituent being determined can be isolated in some weighable form.
Volumetric analysis, on the other hand, derives its name from the fact that the method
used to determine the amount of a constituent involves measuring the volume of a
reagent. Usually, gravimetric analyses involve the following steps:
1- Drying and then accurately weighing representative samples of the material to be
analyzed.
2- Dissolving the samples.
3- Precipitating the constituent in the form of a substance of known composition by
adding a suitable regent.
4- Isolating the precipitate by filtration.
5- Washing the precipitate to free it of contaminants.
6- Drying the precipitate to a constant mass (to obtain an analytically weighable form
of known composition).
7- Calculating the percentage of the desired constituent from the masses of the
sample and precipitate.
Although the techniques of gravimetric analysis are applicable to a large variety of
substances, we have chosen to illustrate them with an analysis that incorporates a number
of other techniques as well. Chloride ion may be quantitatively precipitated from
solution by the addition of silver ion according to the following ionic equation:
Ag+(aq) + Cl–(aq) → AgCl(s)
[1]
Silver chloride is quite insoluble (only about 0.0001 g of AgCl dissolves in 100 mL of
H2O at 20°C); hence, the addition of silver nitrate solution to an aqueous solution
containing chloride ion precipitates AgCl quantitatively. The precipitate can be collected
27
on a filter paper, dried, and weighed. From the mass of the AgCl obtained, the amount of
chloride in the original sample can then be calculated.
This experiment also illustrates the concept of stoichiometry. Stoichiometry is the
determination of the proportions in which chemical elements combine and the mass
relations in any chemical reaction. In this experiment stoichiometry means specifically
the mole ratio of the substances entering into and resulting from the combination of Ag+
and Cl–. In the reaction of Ag+ and Cl– in Equation [1], it can be seen that 1 mol of
chloride ions reacts with 1 mol of silver ions to produce 1 mol of silver chloride.
Example E.1: In a gravimetric chloride analysis it was found that 0.2516 g AgCl was obtained
from an unknown that had a mass of 0.1567 g. What is the % of chloride in the sample?
Solution:
35.45 g Cl
 0.2516 g AgCl  0.06222 g Cl
143.32 g AgCl
% Cl 
0.06222 g Cl
100  39.72 %
0.1567 g sample
Procedure:
1. Obtain an unknown and record its number on your report sheet.
2. On a piece of weighing paper, weigh to the nearest 0.0001 g about 0.2 of your
unknown sample.
3. Transfer the sample quantitatively to a clean 150-ml beaker (do not weigh the
beaker) and label the beaker #1 with pencil. Record the sample weight.
4. Add 100 ml of distilled water and 1 ml of 6 M HNO3 to the beaker.
5. Repeat with sample numbers 2 and 3, and label the beakers #2 and #3,
respectively.
6. Using a different glass rod for each solution, stir until all of the sample has
dissolved. Leave the stirring rods in the beakers. Do not place them on the
desktop.
7. While stirring one of the solutions, add to it about 10 ml of 0.5 M AgNO3 solution.
Place a watch glass over the beaker.
8. Obtain a filter paper (three of these will be needed) and weigh it accurately. (Be
certain that you weigh the paper after it has been folded and torn, not be fore.)
Fold the paper as illustrated in Figure E.1 and fit it into a glass funnel. Be certain
that you open the filter paper in the funnel so that one side has three pieces and
one side has one piece of paper against the funnel – no two pieces on each side.
28
9. Wet the paper with distilled water to hold it in place in the funnel.
10. Completely and quantitatively transfer the precipitate and all warm solution from
the beaker onto the filter, using a rubber policeman and a wash bottle to wash out
the last traces of precipitate. The level of solution in the filter funnel should
always be below the top edge of the filter paper.
11. Wash the precipitate on the filter paper with two or three 5-ml portions of water
from the wash bottle.
12. Finally, pour 5-ml of acetone through the filter. (CAUTION: Acetone is highly
flammable! Keep it away from open flames.)
13. Remove the filter paper, place it on a numbered watch glass, and store it in your
locker until the next period.
14. Repeat the above processes with your other two samples, being sure that you have
numbered your watch glasses so that you can identify the samples.
15. In the next period, when the AgCl is thoroughly dry, weigh the filter papers plus
AgCl and calculate the mass of AgCl. From these data calculate the percentage of
chloride in your original sample.
16. Calculate the standard and relative standard deviations.
The precipitated AgCl must be kept out of bright light, because it is photo- sensitive and
slowly decomposes in the presence of light as follows:
2 AgCl(s)
h
2Ag(s)

+ Cl2(g)
In this equation hν is a symbol for electromagnetic radiation; here it represents radiation
in the visible and ultraviolet regions of the spectrum. This is the reaction used by
Corning to make photosensitive sunglasses.
29
Figure E.1 Filter paper use
30
Name ___________________________________ ID ____________ Sec ___
Experiment 4
Gravimetric Analysis of a Chloride Salt
REPORT SHEET
Trial 1
________
________
________
________
________
Trial 2
________
________
________
________
________
Trial 3
________
________
________
________
________
Percent Cl in original sample
(show calculations)
________
________
________
Average percent chloride
(show calculations)
________
Standard deviation
(show calculations)
________
Relative standard deviations
(show calculations)
________
Mass of sample
Mass of filter paper + AgCl
Mass of filter paper
Mass of AgCl
Mass of Cl in original sample
(show calculations)
Do any of your results differ from the mean by more than two standard deviations?
________
Reported percent chloride
________ ± ________ %
31
EXPERIMENT 5
DETERMINATION OF THE MOLECULAR MASS OF METAL
CARBONATE BY BACK TITRATION
Aim:
The purpose of this experiment is to determine the molecular mass of a metal
carbonate and hence to determine the molar mass of the unknown metal to be identified.
Chemical and other requirements:
Metal Carbonate salt, 1 M Hydrochloric Acid (HCl) , 1 M Sodium Hydroxide.
Phenolphthalein Indicator
Weighing Dished, Spatulas, 250 mL Conical flask , 50 mL graduated pipette, 250 mL
Volumetric flask, Dropper, 10 mL graduated Pipette, Burette, Glass Funnel, Pipette filler.
Introduction:
When a metal carbonate is insoluble in water, some difficulties would arise in titrating
a standard acid against it. These difficulties can be avoided by dissolving a known quantity
of the carbonate in unreacted acid of the acid solution before titrating the unreacted acid with
alkali solution.
The number of moles of HCl that reacted with the base represents the excess of HCl
that has not reacted with the metal carbonate. Therefore, the number of moles of HCl reacted
can be calculated by subtracting the excess number of moles of HCl from the total number of
moles of HCl originally used. Hence, the number of moles of metal carbonate can be found,
so the molar mass is calculated from the original mass of the carbonate. By a simple
calculation, the molar mass of the metal can be determined and compared with that in the
periodic table.
Procedure:
1. Weigh accurately a weighing dish containing not more than (1.5 g) of the metal
carbonate.
2. Empty this carefully into a 250 cm3 conical flask containing 50.0 mL of 1 M
hydrochloric acid and swirl well until no more effervescence is observed.
3. Transfer the mixture into a 250 cm3 volumetric flask and dilute up to the mark using
distilled water and shake well.
4. Pipette 10.0 mL of this solution into a conical flask and titrate against 0.1 M standard
sodium hydroxide. Use phenopthalein as an indicator (2-3drops).
5. Repeat step (4 ) three times.
32
Calculation:
MCO3(s) + 2HCl(aq)  MCl2(aq) + CO2(g) + H2O (l)
1.
2.
3.
4.
5.
6.
7.
8.
9.
Calculate number of moles of NaOH that neutralize 10.0 ml of the solution.
Calculate number of moles of excess HCl in 10.0 mL of the solution.
Calculate number of moles of excess HCl in 250.0 mL solution.
Calculate number of moles of original HCl used.
Calculate number of moles of HCl reacted with metal carbonate.
Calculate number of moles of MCO3 reacted with the acid.
Calculate the molar mass of MCO3.
Calculate the molar mass of Metal.
Identify the metal.
Questions:
1.
Suggest a reason why you can not get an exact molar mass of the metal when compared
to that in the periodical table.
2.
Given the reaction :


BaCO3(s) 
BaO(s) + CO2(g)
If 85.6 g of impure sample BaCO3 is used, then 10.8 g of CO2 is obtained. What is the
percentage of BaCO3 in the sample.
3.
Given the reaction:
(NH4)2CO3(aq) + FeCl2(aq)  FeCO3(s) + 2NH4Cl(aq)
What mass of FeCO3(s) is obtained if 45.0 mL of 0.80 M of (NH4)2CO3 reacted with
excess of FeCl2.
Suggested table for your results.
Mass of metal carbonate (g)
NaOH
Titrant
Volume of analyte
HCl
Analyte
Indicator
Titrations
Readings of the burette
Initial (I) mL
Final (F) mL
1
2
3
4
Average volume
33
10.0 mL
Phenolphthalein
Volume used = F-I
(mL)
Experiment 6
Spectrophotometric Analysis of a Commercial Aspirin Tablet
Objective:
To determine the % of active ingredient chemical of a commercial aspirin.
Materials and chemicals:
UV/VIS spectrophotometer and polystyrene cuvettes
Hot plate
4 15 mL test tubes
2 10 mL graduated pipette
1 5 mL graduated pipette
1 150 mL beaker (or 250 mL beaker)
1 10 mL graduated cylinder
1 100 mL volumetric flask
1 50 mL volumetric flask
1 watch glass
1 commercial aspirin tablet
50 mL ~ 0.000725 M salicylic acid (reagent grade)
50 mL 0.02 M FeCl3 (buffered to pH = 1.6 with HCl/KCl)
10 mL 1.0 M NaOH
Introduction:
Acetyl salicylic acid (ASA) is one of the oldest synthetic drugs. First synthesized in
Germany by the Bayer company and marketed under the name “Aspirin” it has remained
one of the most popular “over the counter” drugs of all time. Its main effect is as a pain
killer and fever depressant, but in addition there is strong evidence that in low daily
dosages it lowers the incidence of heart attacks. In the last few decades other drugs such
as acetaminophen (commercial trade name Tylenol) and ibuprofen (trade name Advil)
have taken much of the market for ASA, but ASA remains an important and widely used
medicine.
Drugs, in addition to their active compound, often contain other inactive ingredients
(called excipients in the pharmaceutical industry) such as binders, fillers, dyes, drying
34
agents, etc. Te content of active ingredient in a tablet will always be stated on the
package. In this experiment we will determine the percent active compound in a
commercial aspirin tablet. Aspirin is the trade name for acetylsalicylic acid (ASA). The
ASA in the tablet will be reacted with Fe3+, forming an intensely violet coloured
complex. The concentration of the complex will be determined by means of
spectrophotometry, using a UV/VIS spectrophotometer. Finally, we will be able to
calculate the weight and the weight% of ASA in the commercial tablet.
Procedure:
A) Preparation of the Fe-salicylate standard solutions and the calibration
curve
1. Use a 10 mL graduated pipette to transfer 1.0, 3.5, 6.5, and 9.0 mL of 0.000725 M
salicylic acid solution into four separate test tubes. Label these solutions SA1,
SA2, SA3, SA4, respectively.
2. Make the total volume in each test tube equals to 10.0 mL by adding 0.02 M FeCl3
solution using a 10 mL graduated pipette.
3. Measure the absorbance of the four standard solutions at 530 nm using the cuvette
provided. Start with the lowest concentration (solution SA1). Rinse the cuvette
with the next solution (SA2) then measure the absorbance of this solution, etc.
B) Analysis of the ASA in a commercial aspirin tablet
(This procedure can be done at the same time as preparing the standard solutions)
1. Weigh a commercial aspirin tablet directly into 150 ml beaker.
2. Add 3.5 mL of 1.0 M NaOH (use 10 ml graduated cylinder).
3. Place the beaker on a hot plate at the lowest temperature. Slowly heat the mixture
for about 5 minutes or until all solid is dissolved (cover with watch glass to
prevent spattering). Then allow the solution to cool and the undissolved material
to settle down.
4. Decant the liquid solution into a 100 mL volumetric flask and fill to the mark with
distilled water. Label this solution “STOCK ASA UNKNOWN”.
5. Transfer 2.0 mL of STOCK ASA UNKNOWN solution (with a 5 mL graduated
pipette) to 50 mL volumetric flask. Fill the flask to the mark with 0.02 M FeCl3
solution. Label this solution “ASA UNKNOWN”. The presence of binders and
fillers in the tablet may make the solution cloudy, but this will disappear when
diluted with the acidified Fe3+ solution.
6. Measure the absorbance of the ASA UNKNOWN solution at 530 nm.
35
Results and Calculations:
A) Standard solutions and Beer’s law calibration curve
1- Fill in the following table:
Solution
SA1
SA2
SA3
SA4
Absorbance
Concentration, M
2- Plot a graph of Absorbance vs. concentration C (mol/L)
B) Analysis of unknown aspirin tablet
1- Mass of aspirin tablet = _________________________ g.
2- Absorbance of “ASA UNKNOWN” = _______________________
3- Concentration of ASA UNKNOWN as determined from the calibration
curve,
CASA UNKNOWN = ______________________________ mol/L
4- Concentration of “STOCK ASA UNKNOWN”,
CSTOCK ASA UNKNOWN = (50.0/2.00) × CASA UNKNOWN =
______________________________________________________
mol/L
5- Moles of ASA in 100 mL solution = __________________________________________ mol
6- Mass of ASA = mol × molar mass =
__________________________________________
7- % of ASA in aspirin = _______________________________________________ %
36
g
Experiment 7
Chemical Reactions
Objective:
To observe some typical oxidation-reduction reactions and metathesis reactions,
identify some of the products and become familiar with writing balanced chemical
equations, including net ionic equations.
Introduction:
Chemical equations represent what happens in a chemical reaction. Before an equation
can be written for a reaction, someone must establish what the products are. Products are
identified by their chemical and physical properties as well as by analyses.
In this experiment, you will observe that in some cases gases are produced,
precipitates are formed, or color changes occur during the reactions. These are all
indications that a chemical reaction has occurred. To identify some of the products of the
reactions, consult Table A.1, which lists some of the properties of the substances that
could be formed in these reactions.
Table A.1 Properties of Reaction Products
Water-soluble solids
Water-insoluble solids
Manganese oxyanions
Gases
KCl: white (colorless
solution)
NH4Cl: white
(colorless solution)
KMnO4: purple
MnCl2: pink (very
pale)
Cu(NO3)2: blue
CuS: very dark blue
or black
Cu2S: black
BaCrO4: yellow
BaCO3: white
PbCl2: white
MnO2: black or
brown
MnO4–: purple
MnO42–: dark green
MnO43–: dark blue
H2: colorless; odorless
NO2: brown; pungent
odor (TOXIC)
NO: colorless slight,
pleasant odor
CO2: colorless;
odorless
Cl2: pale yellow-green;
pungent odor
(TOXIC)
SO2: colorless; choking
odor (as from
matches ) (TOXIC)
H2S: colorless; rottenegg odor (TOXIC)
37
I. Oxidation-Reduction Reactions
Many metals react with acids to liberate hydrogen and form the metal salt of the acid.
The noble metals do not react with acids to produce hydrogen. Some of the unreactive
metals do react with nitric acid, HNO3; however, in these cases gases that are the oxides
of nitrogen are formed rather than hydrogen.
II. Metathesis Reactions
In molecular equations for many aqueous reactions, cations and anions appear to
exchange partners. These reactions conform to the following general equation:
AX + BY  AY + BX
[1]
Such reactions are known as metathesis reactions. For a metathesis reaction to lead to a
net change in solution, ions must be removed from the solution. In general, three
chemical processes can lead to the removal of ions from solution, thus serving as a
driving force for metathesis to occur:
1. The formation of a precipitate.
2. The formation of a weak electrolyte or nonelectrolyte.
3. The formation of a gas that escapes from solution
A. Formation of a Precipitate
The reaction of barium chloride with silver nitrate is a typical example:
BaCl2(aq) + 2AgNO3(aq) → Ba(NO3)2(aq) + 2AgCl(s)
[2]
This form of the equation for this reaction is referred to as the molecular equation.
Because we know that the salts BaCl2, AgNO3, and Ba(NO3)2 are strong electrolytes and
are completely dissociated in solution, we can more realistically. write the equation as
follows:
Ba2+(aq) + 2Cl– (aq) + 2Ag+(aq) + 2NO3– (aq) → Ba2+(aq) + 2NO3– (aq) + AgCl(s)
[3]
This form, in which all ions are shown, is known as the complete ionic equation.
Reaction [2] occurs because the insoluble substance AgCl precipitates out of solution.
38
The other product, barium nitrate, is soluble in water and remains in solution. We see that
Ba2+ and NO3– ions appear on both sides of the equation and thus do not enter into the
reaction. Such ions are called spectator ions. If we eliminate or omit them from both
sides, we obtain the net ionic equation:
Ag+(aq) + Cl– (aq) → AgCl(s)
[4]
This equation focuses our attention on the salient feature of the reaction: the formation
of the precipitate AgCl. It tells us that solutions of any soluble Ag+ salt and any soluble
Cl– salt, when mixed, will form insoluble AgCl. When writing net ionic equations,
remember that only strong electrolytes are written in the ionic form. Solids, gases,
nonelectrolytes, and weak electrolytes are written in the molecular form.
B. Formation of a gas
Molecular equation:
2HCl(aq) + Na2S(aq) → 2NaCl(aq) + H2S(g)
Complete ionic equation:
2H+(aq) + 2Cl– (aq) + 2Na+(aq) + S2– (aq) → 2Na+(aq) + 2Cl– (aq) + H2S(g)
[5]
[6]
Net ionic equation:
2H+(aq) + S2– (aq) → H2S(g)
[7]
C. Formation of a Weak Electrolyte
Molecular equation:
HNO3(aq) + NaOH(aq) → H2O(ℓ) + NaNO3(aq)
[8]
Complete ionic equation:
H+(aq) + NO3– (aq) + Na+(aq) + OH– (aq) → H2O(ℓ) + Na+(aq) + NO3– (aq)
[9]
Net ionic equation:
H+(aq) + OH– (aq) → H2O(ℓ)
39
[10]
In order to decide if a reaction occurs, we need to be able to determine whether or not a
precipitate, a gas, a nonelectrolyte, or a weak electrolyte will be formed. The following
brief discussion is intended to aid you in this regard. Table A.2 summarizes solubility
rules and should be consulted while performing this experiment.
Table A.2 Solubility Rules
Water-soluble salts
Na+, K+, NH4+
All sodium, potassium, and ammonium salts are soluble.
NO3–, ClO3–, C2H3O2–
All nitrates, chlorates, and acetates are soluble.
Cl–
All chlorides are soluble except AgCl, Hg2Cl2, and PbCl2*.
Br–
All bromides are soluble except AgBr, Hg2Br2, PbBr2,* and
HgBr2*.
All iodides are soluble except AgI, Hg2I2, PbI2, and HgI2.
All sulfates are soluble except CaSO4,* SrSO4, BaSO4, Hg2SO4
and PbSO4.
I–
SO42–
Water-insoluble salts
CO32–, SO32–, PO43–, CrO42–
OH–
S2–
All carbonates, sulfites, phosphates, and chromates are insoluble
except those of alkali metals and NH4+.
All hydroxides are insoluble except those of alkali metals and
Ca(OH)2,* Sr(OH)2,* and Ba(OH)2.
All sulfides are insoluble except those of the alkali metals, alkaline
earths, and NH4+.
*Slightly soluble.
The common gases are CO2, SO2, H2S, and NH3. Carbon dioxide and sulfur dioxide may
be regarded as resulting from the decomposition of their corresponding weak acids,
which are initially formed when carbonate and sulfite salts are treated with acid:
H2CO3(aq) → H2O(ℓ) + CO2(g)
[11]
and
H2SO3(aq) → H2O(ℓ) + SO2(g)
[12]
Ammonium salts form NH3 when they are treated with strong bases:
NH4+(aq) + OH– (aq) → NH3(g) + H2O(ℓ)
[13]
Which are the weak electrolytes? The easiest way of answering this question is to identify
all of the strong electrolytes, and if the substance does not fall in that category it is a
weak electrolyte. Note, water is a nonelectrolyte. Strong electrolytes are summarized in
Table A.3.
40
In the first part of this experiment, you will study some oxidation-reduction reactions
and in the second part you will study some metathesis reactions. In some
instances it will be very evident that a reaction has occurred, whereas in others it will not
be so apparent. In the doubtful case, use the guidelines above to decide whether or not a
reaction has taken place. You will be given the names of the compounds to use but not
their formulas. This is being done deliberately to give practice in writing formulas from
names.
Table A.3 Strong Electrolytes
Salts
All common soluble salts
Acids
HClO4, HCl, HBr, HI, HNO3, and H2SO4 are strong electrolytes; all
others are weak.
Bases
Alkali metal hydroxides, Ca(OH)2, Sr(OH)2, and Ba(OH)2 are strong
electrolytes; all others are weak.
41
Procedure:
I. Oxidation-Reduction Reactions
Add a small piece of zinc to a test tube containing 2 mL of 2 M HCl, and note what
happens.
1. Record your observations
2. Suggest possible products for the observed reaction: Zn(s) + HCl(aq)  ?
Place a 1-in. piece of copper wire in a clean test tube. Add 2 mL of 2 M HCl, and note if
a reaction occurs.
3. Record your observations
4. Is copper an active or an inactive metal?
WHILE HOLDING A CLEAN TEST TUBE IN THE HOOD, place a 1-in. piece of
copper wire in it and add 1 mL of 6M acid, HNO3.
5.
6.
7.
8.
Record your observations.
Is the gas colored?
Suggest a formula for the gas.
Based on the color of the solution, what substance is present in the solution?
Potassium permanganate, KMnO4, is an excellent oxidizing agent in acidic media. The
permanganate ion is purple and is reduced to the manganous ion, Mn2+, which has a very
faint pink color.
Place I mL of 0.1 M sodium oxalate, Na2C2O4, in clean test tube. Add 10 drops of 4 M
sulfuric acid. Mix thoroughly. To the resulting solution, add 10 drops of 0.02 M KMnO4
and stir. If there is no obvious indication that a reaction has occurred, warm the test tube
gently in a hot water bath.
9. Record your observations. Was KMnO4 reduced to Mn2+?
II. Metathesis Reactions
The report sheet lists 8 pairs of chemicals that are to be mixed. Use about 15 drops of the
reagents to be combined as indicated on the report sheet. Mix the solutions in small test
tubes and record your observations on the report sheet. If there is no reaction, write N.R.
42
Name______________________________________ ID _____________ Sec ___
Experiment 7:
Chemical Reactions
REPORT SHEET
I. Oxidation–Reduction Reactions
1.
2. Zn(s) + HCl(aq) 
3.
4.
5.
6.
7.
8.
9.
II. Metathesis Reactions
10. Copper(II) sulfate + sodium carbonate
Observations
Molecular equation
Net ionic equation
11. Copper(II) sulfate + barium chloride
Observations
Molecular equation
Net ionic equation
12. Sodium carbonate + hydrochloric acid
Observations
Molecular equation
Net ionic equation
13. Nickel chloride + sodium carbonate
43
Observations
Molecular equation
Net ionic equation
14. Ammonium chloride + sodium hydroxide
Observations
Molecular equation
Net ionic equation
15. Sodium acetate + hydrochloric acid
Observations
Molecular equation
Net ionic equation
16. Potassium chloride + sodium nitrate
Observations
Molecular equation
Net ionic equation
44
EXPERIMENT 8
DETERMINING THE MOLAR MASS OF A GAS
Aim:
The purpose of this experiment is to determine the molar mass and density of a gas
using ideal gas law.
Introduction:
In the gas phase, all substances show similar physical behavior. Volumes of
different gases respond in almost exactly the same way to changes in amount, in pressure
or in temperature.
Gas Laws
Boyle’s Law (1662)
At a given temperature and number of moles, the pressure and volume of a gas are
inversely proportional in the form of an equation:
PV  constant
Charles’s Law (1787)
At a given pressure and number of moles, the volume and temperature of a gas are
directly proportional in the form of an equation:
V
 constant
T
45
Avogadro’s Hypothesis (1811)
At a given temperature and pressure, the volume and mass (number of moles, n) of
a gas are directly proportional in the form of an equation:
V
 constant
n
Ideal Gas Law
Combining the relationships above yields the Ideal Gas Law.
PV  nRT
Each quantity in the equation is usually expressed in the following units:
P = Pressure, measured in atmospheres.
V = Volume, measured in liters.
n = Amount of gas, measured in moles.
T = Absolute temperature, measured in Kelvin.
R = The ideal gas constant, which has a value of 0.0821 L atm/mol K.
At low pressure ( ~ 2-3 atmosphere or below) and high temperature (greater than
0C), most gases obey the ideal gas equation.
It is convenient to rewrite the Ideal Gas Law in a slightly different form,
expressing the amount in grams rather than moles.
n
Substituting for n in the Ideal Gas Law,
Solving for MM
MM 
or solving for density d 
m
MM
PV 
mRT
PV
m PMM

V
RT
46
mRT
MM
Procedure:
1. Get instructions from your instructor on how to operate the gas cylinder – there
will be some valves you must not touch.
2. Weigh a dry volumetric flask together with its stopper to the nearest 0.001g.
Record the mass on your data sheet.
3. Remove the stopper, insert the glass delivery tube from the gas cylinder or
generator so that it reaches the bottom of the flask, and open the valve so that the
gas passes through for at least one minute. Keep the flask upright throughout.
4. Slowly remove the delivery tube and quickly close the flask with the stopper.
5. Weigh the flask with the stopper again to the nearest 0.001g. Preferably, using the
same balance used in step 2 throughout the experiment.
6. Repeat steps 3,4 and 5 and check that there is no further change in mass. If this is
not the case, repeat these steps again, if necessary, until the mass is constant.
7. Fill the flask with water and insert the stopper, so that the excess water is pushed
out. Dry the outside of the flask and weigh it, full of water, on a balance to the
nearest 0.1g.
8. Note room temperature and atmospheric pressure.
The following table gives values for the density of the air under various conditions of
temperature and pressure. If your conditions do not correspond to any of those quoted,
you should estimate the appropriate value.
Table 1: Density of air (g cm-3) at different temperatures and pressures.
740 mmHg
750 mmHg
760 mmHg
770 mmHg
780 mmHg
15ºC
0.00119
0.00121
0.00123
0.00124
0.00126
17ºC
0.00119
0.00120
0.00122
0.00123
0.00125
19ºC
0.00118
0.00119
0.00121
0.00123
0.00124
Density of water = 1.00 g cm-3
47
21ºC
0.00117
0.00119
0.00122
0.00122
0.00123
23ºC
0.00116
0.00118
0.00119
0.00121
0.00122
25ºC
0.00115
0.00117
0.00119
0.00120
0.00122
Calculation:
1. Calculate the molar mass of the gas at room temperature and pressure.
2. Calculate the density of gas at the experimental conditions and compare it with the
density at STP.
Questions:
1. What value does the experiment give for the molecular mass of the gas? Identify
the gas.
2. In Step (4) why were you told to remove the delivery tube slowly?
3. Why is a less accurate balance adequate for weighing the flask full of water?
48
Name______________________________________ ID _____________ Sec ___
Experiment 8:
DETERMINING THE MOLAR MASS OF A GAS
DATA SHEET
Mass of flask filled with air (g)
---------------------------------------------
Mass of flask filled with gas (g)
---------------------------------------------
Mass of flask filled with water (g)
--------------------------------------------
Room temperature (K)
--------------------------------------------
Atmospheric pressure (atm)
--------------------------------------------
Density of air under conditions of experiment (g cm-3) ----------------------------------------
Mass of water (g)
--------------------------------------------
Volume of water = Volume of flask (mL)
--------------------------------------------
Mass of gas (g)
---------------------------------------------
The molar mass of the gas (g/mole)
---------------------------------------------
49
Experiment 9
Separation of a Mixture by Paper Chromatography
Objective:
To separate, detect, and identify the components of a mixture by using paper
chromatography.
Introduction:
Chromatography is the study of separations of mixtures and is often used to identify
unknown components in mixtures. Paper chromatography, which is used here, is just one
of several chromatographic techniques available.
In paper chromatography, a drop of solution containing a mixture of substances is
placed near one end of a rectangular piece of filter paper (the stationary phase). The end
of the paper is immersed in a liquid (the mobile phase) to a point that is just below the
spot where the drop was placed on the paper. Capillary action causes the liquid to flow up
the filter paper. When the liquid reaches the spot, the components of the mixture will
begin to migrate upward the mobile phase. Each component will have a characteristic
chemical affinity for the paper and a characteristic chemical affinity for the liquid. A
component with a strong affinity for the paper and a weak affinity for the liquid will
move more slowly than a component with a weaker affinity for the paper and a stronger
affinity for the liquid. Because each component of mixture has its own characteristic
affinities, each component will travel up the paper at its own characteristic rate. If the
paper is sufficiently large, all the components can be separated by the time the liquid has
reached the top of the paper.
Each component will now paper as separate spot. If the components are highly
colored, the spots will be visible. You can convert weakly colored or colorless spots to
highly colored spots by spraying them with substances that react with the components in
the spots. The filter paper will now contain a vertical array of colored spots arranged
according to their characteristic rates of ascent.
The distance traveled by a component of a spot with respect to the distance traveled
by the pure liquid is a measure of that component's competitive affinities for the
stationary and mobile phases. We define the component's Rf (retention factor) value I
those terms:
Rf 
distance traveled by spot
distance traveled by liquid
50
The Rf value of a compound is a characteristic of the compound, the support, and then
solvent used, and it serves to identify the constituents of a mixture. An estimate of the
relative amount of each of the constituents in a mixture can be made from the relative
intensities and sizes of the various bands in a chromatogram.
You will examine the paper chromatography of Fe3+, Cu2+, and Ni2+. Fe3+ in water
imparts a red-brown or rust color and thus will produce a rust-colored band on the paper.
Although Cu2+ is blue, the color is faint and not easily detected, especially if copper is
present in small amounts. In aqueous solution, however, Cu2+ reacts with NH3 (from
ammonium hydroxide) to form a complex ion, [Cu(NH3)4]2+, which is deep blue and
therefore readily observed. Finally, Ni2+ reacts with an organic reagent,
dimethylglyoxime, to produce strawberry-red color. The mobile phase will be a mixture
of 90% acetone and 10% 6 M HCl.
You will be able to determine the Rf value for each cations by observing its ascent in
the absence of the other cations. You will also subject a known mixture of all the cations
to chromatography so that you can see that the same Rf values are obtained with a
mixture. The components in an unknown mixture containing one or more of the cations
being studied will also be identified on the basis of its chromatography and derived Rf
values. Finally, a mixture containing Fe3+, Ni2+, and a trace contaminant of Cu2+ will be
examined in order to determine the relative amounts of components of a mixture.
There are two simple ways of performing this experiment: either by circular
horizontal chromatography or by ascending chromatography using paper strips. The latter
technique will be used in this experiment.
Procedure:
1. Obtain a piece of chromatography filter paper and use a pencil (not a pen) to draw
a line 2 cm from one of the longer edges. Mark the line with a pencil dot at six
places which are equally spaced. Place the filter paper on top of another larger
piece of filter paper or a paper towel, which will serve as an absorbing pad.
2. Fill your capillary pipet by dipping the end into the Cu2+solution. Allow the
solution to rise by capillary action.
51
3. Withdraw the pipet and touch the inside of the vial with the tip of the pipet to
remove the hanging drop. Spot the filter paper on the first pencil dot by touching it
with capillary held perpendicular to the paper. Allow the solution to flow out of
the capillary until a spot having about 3-mm diameter is obtained. Dry the filter
paper completely by waving it in the air.
4. In the same way, apply a single drop of Ni2+, and then Fe3+ solution to the
subsequent pencil dots.
5. Using a pencil, write the symbol of the cations close to the corresponding spot.
6. Spot the remaining dots with 1 drop of the known, the unknown, and the trace
Cu2+ solutions.
7. Pour the developing solvent to a depth of 1 cm (about 20 mL) in the bottom of a
600-mL beaker.
8. Fold the filter paper and staple to make a cylinder shape as shown in the figure
below and place it in the developing chamber.The 2-cm line should be above the
surface of the liquid. Be certain that the spots on the paper are completely dry
before you place the paper in the beaker. Make sure that the paper does not touch
the wall of the beaker and that the bottom of the paper is resting in the solution.
9. Cover the beaker carefully with a watch glass. Do not disturb the beaker while the
chromatograms are developing.
10. When the solvent front has nearly reached the top of the paper, carefully remove
the watch glass and the paper. Place the wet paper on the paper towel and
immediately mark the solvent front with a pencil. Allow the filter paper to dry by
fanning the air with it.
11. What are the colors of the spots that you are able to see (1)? Which of the known
ions do you detect without restoring to the use of any other development (2)?
Developing reactions
1. IN THE HOOD, pour about 5 mL of 15 M ammonia into a clean, shallow dish and
rest the filter paper on the top of the dish. Do not permit the paper to dip into the
solution. What color develops (3)? What ions dos this indicate (4)?
2. To detect the third ion, dip a new piece of filter paper into a solution of
dimethylglyoxime; then, using the new piece as a brush, paint the test filter paper.
What color develops (5)? What ions does this indicate (6)?
52
3. Measure the vertical distance that the approximate center of each of these spots
and the pure solvent has traveled from the original mark on the 2-cm line. Record
your results (7). Calculate the Rf value for each cation (8).
4. Write the letter label of your unknown (9), measure the distance to each of the ring
fronts in your unknown (10) and calculate the Rf values (11). What ions are
present in your unknown (12)?
5. Record your observations on the relative amounts of the ions based on color
intensities in the solution with trace Cu2+ (13).
53
Name______________________________________ ID _____________ Sec ___
Experiment 9:
Separation of a Mixture by Paper Chromatography
REPORT SHEET
1. Color of spots: Cu2+___________ Ni2+____________ Fe3+___________
2. Ion requiring no development ___________
3. Ammonia develops a ___________ color.
4. The ion that causes the color of (3) is _________
5. Dimethylglyoxime develops a ______________ color.
6. The ion that causes the color of (5) is __________
7. Ring front distances for knowns are: solvent: _______ mm;
Cu2+ _________ mm; Ni2+ __________ mm ; Fe3+ _________ mm.
8. Rf values for knowns: Cu2+ ________ ; Ni2+ ________ ; Fe3+ ________ .
9. Unknown label ________
10. Ring front distances for unknowns:
__________ mm ; __________ mm ; __________ mm.
11. Rf values for unknowns:
____________ mm ; ____________ mm ; ____________ mm ; .
12. Ions present in unknown __________________
13. Observation on relative amounts _________________________________
______________________
54
55