Chem 26.1 Lab Manual 2017 Edition (2019) PDF
Chem 26.1 Lab Manual 2017 Edition (2019) PDF
Chem 26.1 Lab Manual 2017 Edition (2019) PDF
College of Science
University of the Philippines, Diliman, Quezon City
CHEMISTRY 26.1
Introduction to Quantitative Chemical Analysis
LABORATORY MANUAL
2017 EDITION
Chemistry 26.1: Introduction to Quantitative Chemical Analysis Laboratory Manual (2017 Edition)
Foreword
This laboratory manual is intended for use in Analytical Chemistry Laboratory taken by non-
Chemistry major students. This manual was developed, researched and revised by the Analytical
Chemistry Academic Group of the Institute of Chemistry. The main objective in this revised edition
is to provide students with the best practical procedures for learning chemistry by incorporating and
amplifying features that enhance their understanding of basic analytical techniques used in chemical
measurements. In addition, this manual was also developed to provide guidance in the area of
general laboratory safety. It is a part of our overall effort to establish basic, safe operating practices
so that students and teachers can do effective teaching and research programs in a safe and healthy
environment.
I would like to acknowledge the people that have made considerable contributions to the completion
of this current edition of Chemistry 26.1 laboratory manual.
Acknowledgement
Following the 2009 and 2013 revision, the Analytical Chemistry Academic Group has produced this
third version of the new laboratory manual, set to be released in 2017, for the use of students of the
University of the Philippines Diliman enrolled in Chemistry 26.1: Analytical Chemistry Laboratory.
The 2017 edition adapts most of the content and revisions from the 2013 edition. What differentiates
this 2017 edition from the previous ones are as follows: (1) modified procedures for most experiments,
following microscaling and waste disposal considerations, and (2) modified data sheets, both
student’s and instructor’s copies, for most experiments.
This 2017 version is a collective effort of the following instructors who patiently reviewed and revised
the experiments and details of the laboratory manual: Ms. Rosemarie Elloisa Acero, Ms. Joyce Lyn
Garcia and Mr. Cris Angelo Pagtalunan.
The cover of the 2017 edition is by Rajelle Hernandez, a BS Chemistry graduate of UP Diliman.
Contents
Students’ Guide ................................................................................................................................................. 1
Students’ Guide
Pre-Laboratory Discussion
The instructor will give a pre-laboratory discussion for each experiment. This is composed of a short
introduction on the type of analysis to be performed and the procedures of the experiment.
Post-Laboratory Discussion
After each experiment has been performed, the instructor will give a post-laboratory discussion. This is
composed of a detailed discussion of the concepts, results of the experiment and possible sources of error, as
well as proper calculations and solutions.
Laboratory Performance
The laboratory performance of the student is evaluated for every experiment based on the standards set by the
instructor. The student is evaluated through attendance during experiment day and through the laboratory
procedures and techniques s/he performs during the experiment.
Laboratory Group
Each student will work on each experiment with a partner or groupmates (in case of odd-numbered classes).
Locker
Each pair or group will be assigned a locker, which should be kept clean and secure at all times. The locker
should also be lined with newspaper, manila paper or any wrapper to avoid moisture accumulation and
spillage. The contents of each locker will be checked-out during the start of the semester and will be checked-
in at the last day of classes, to account for any loss which might have occurred during the semester.
Data Notebook
The data notebook is required for each student. It should be an 8.5” x 11” bound notebook with 50-60 leaves.
All the right-hand pages of the notebooks should be numbered at the upper right hand corner. The left-hand
pages only serve as scratch and area for note-taking. The first page/s of the notebook will serve as the table of
contents, which should be updated regularly. This page should have the following format:
Figure SG-1. The format of the table of contents of the laboratory notebook.
Each student is required to submit a pre-laboratory report for each experiment a day prior the actual
experiment. This report includes the following: (a) the OBJECTIVES of the experiment, (b) a LIST OF ALL
GLASSWARE, MATERIALS, EQUIPMENT AND REAGENTS (quantity included) to be used in the experiment,
(c) SCHEMATIC DIAGRAM OF THE PROCEDURE, (d) DIAGRAM/S OF SPECIAL SETUP/S, (e) WASTE
DISPOSAL procedures and (f) CALCULATIONS for solution preparation.
The data notebook will serve as the student’s guide during the experiment since all laboratory manuals will not
be allowed to be open inside the laboratory. The calculations for solution preparation should include the
volume of solution required for dilution, concentration of a solution, mass of substances to be weighed, etc. for
all solutions to be prepared.
Attendance
A student is considered late if s/he comes to class 15-30 minutes after the class has started.
A student is considered absent if: a) s/he does not have complete PPE and/or b) s/he comes to class 30 minutes
after the class has started. An absence, whether excused or unexcused, merits a student a grade of 0 for
laboratory performance. No make-up experiments are allowed for any type of absences. For excused absences,
however, the student will still be allowed to submit the required report for the experiment.
Monitors
Monitors will be assigned for each experiment. It is the duty of the monitor to: (a) borrow and return all floating
glassware and equipment needed by the class for each experiment, (b) lead and manage solution preparation
for the whole class, (c) maintain proper decorum and cleanliness of the class during the experiment, and (d)
check if the laboratory room is in order and all lockers are secured at the end of the class.
Overtime
Overtime during experiments is not tolerated.
Data Sheets
The data sheets are found at the end of the laboratory manual. All needed data must be presented after the
experiment. The instructor’s copy of the data sheet should be submitted to the instructor at the end of each
experiment. It will be checked by the instructor at the end of the period. On the other hand, the student’s copy
of the data sheet will be submitted along with the calculated results and report (ATQ Report or Formal Report).
Only blue or black pens must be used for filling the data sheet. Pencils and liquid erasers are not allowed. If an
entry is found to be wrong, it should be struck-through and signed, before being replaced by the correct data.
Post-Laboratory Reports
Each student must submit the completed data sheet and a laboratory report (either an ATQ or an FR) ONE
WEEK AFTER THE LAST DAY OF THE EXPERIMENT BEFORE THE CLASS STARTS. Late submission gives
the student a grade of 0 for the ATQ/FR of that experiment. Non-submission of a report gives the student a
grade of INC, if the total standing of the student is passing.
As guide for students in correct scientific paper writing, a trial FR will be submitted. This will serve as training
for students in writing an FR. However, this trial FR will be graded as an ATQ report.
Academic Dishonesty
Academic dishonesty (copying in examination, plagiarism, among others) is not tolerated. A student caught to
be performing academic dishonesty acts shall be given a grade of 5.0 and may be subjected to disciplinary action
by the University.
Plagiarism
(Reference: University of Washington Psychology Writing Center. 1997. Plagiarism and Student Writing. Web.
2013. <http://www.psych.uw.edu/writingcenter/writingguides/pdf/plag.pdf>)
Plagiarism occurs when one uses the ideas or writings of another as his/her own without giving due credit. A
student commits plagiarism by any of the following:
a) using another writer’s words without proper citation,
b) using another writer’s ideas without proper citation,
c) citing your source but reproducing the exact words of a printed source without quotation marks,
d) borrowing the structure of another author’s phrases or sentences without crediting the author who
wrote it,
e) borrowing all or part of another student’s paper, or using someone else’s outline to write your own
paper, and
f) using a paper writing service or having a friend write the paper for you.
D. Working Equations
E. Sample Calculations
Other remarks:
• Maximum of four (4) pages (including references but excluding appendix)
• Appendix should still be paginated. A two column format should be followed in the body until references. Appendix
should follow single column format.
• Attach your data sheets when submitting your formal report
Chemistry 26.1: Introduction to Quantitative Chemical Analysis Laboratory Manual (2017 Edition)
EXPERIMENT 1
Application of Statistical Concepts in the Determination of
Weight Variation in Samples
OBJECTIVES
At the end of the experiment, the student should be able to:
1) use an analytical balance properly;
2) gain an understanding of some concepts of statistical analysis; and
3) apply statistical concepts in analytical chemistry.
INTRODUCTION
In dealing with the numerical results in an experiment, it is important to assess both the accuracy and precision
of these data. Accuracy refers to the closeness of a measurement to the true or accepted value. Precision refers
to the closeness of the measurements that have been obtained using the same method. The accuracy and
precision of measurements are evaluated using statistical tests.
1. Mean
One of the most common measures of central tendency is the mean (or average).
∑𝑛𝑖=1 𝑋𝑖 (𝑋1 + 𝑋2 + 𝑋3 + ⋯ 𝑋𝑛 )
𝑋̅ = = (1.1)
𝑛 𝑛
2. Median
The median is the middle value in a set of data that has been arranged in increasing or decreasing order. The
median is useful when a set of data contains an outlier, a result that differs significantly from the rest of the data
in the set. For an odd number of results, the median can be evaluated directly. For an even number, the average
of the middle pair is used.
Measures of Accuracy
Accuracy is expressed in terms of absolute error or relative error.
1. Absolute Error, E
Absolute error is the difference between the experimental and true value. Xi is the experimental value and Xt is
the true value.
𝐸 = 𝑋𝑖 − 𝑋𝑡 (1.2)
2. Relative Error, Er
Relative error is the absolute error divided by the true value. Relative error is usually expressed in percent.
𝑋𝑖 − 𝑋𝑡
𝐸𝑟 = × 100% (1.3)
𝑋𝑡
Measures of Precision
1. Variance, s2
Variance is a more statistically useful measure of precision.
∑𝑛𝑖=1(𝑋𝑖 − 𝑋̅)2
𝑠2 = (1.4)
𝑛−1
2. Standard Deviation, s
Standard deviation is just the square root of the variance.
∑𝑛 (𝑋𝑖 − 𝑋̅)2
𝑠 = √ 𝑖=1 (1.5)
𝑛−1
Note that for large values of n, the number of degrees of freedom, (n – 1), approaches n. Consequently, the
sample standard deviation approaches the population standard deviation.
∑𝑛𝑖=1(𝑋𝑖 − 𝜇)2
𝜎 =√ (1.6)
𝑛
4. Coefficient of Variation, CV
When the relative standard deviation is in percent, it is called the coefficient of variation.
𝑠
𝐶𝑉 = × 100 % (1.8)
𝑋̅
𝑛1 𝑛2 2 𝑛3
∑𝑖=1 (𝑋𝑖 − ̅̅̅
𝑋1 )2 + ∑𝑗=1 (𝑋𝑗 − ̅̅̅
𝑋2 ) + ∑𝑘=1 ̅̅̅3 )2
(𝑋𝑘 − 𝑋
𝑠𝑝𝑜𝑜𝑙𝑒𝑑 = √ (1.9)
𝑛1 + 𝑛2 + 𝑛3 +. . . −𝑛𝑠
where n1 is the number of data in set 1, n2 is the number of data in set 2, and so forth. The term ns is the number
of data sets that are being pooled.
6. Range, R
Range is the difference between the highest and lowest value in a set of measurements.
7. Relative Range, RR
Range may also be expressed in relative terms.
𝑅
𝑅𝑅 = × 1000 𝑝𝑝𝑡 (1.11)
𝑋̅
Confidence Interval, CI
Confidence interval provides a range of values within which the population mean is expected to lie at a specified
confidence level. The boundaries of the confidence interval are called confidence limits.
𝑡𝑠
𝐶𝐼 = 𝑋̅ ± (1.12)
√𝑛
t is dependent on the confidence level and degrees of freedom. The 95% level which incorporates about two
standard deviation units is often used in getting the confidence interval.
Grubbs’ Test
An objective criterion must be used in rejecting suspected outliers. Grubbs’ test is one of the most commonly
used for detecting outliers. This test is applicable only for data sets containing one suspected outlier. The
experimental value, g, is calculated and compared with tabulated critical g values.
|𝑋𝑖 − 𝑋̅|
𝐺𝑐𝑎𝑙𝑐 = (1.13)
𝑠
Xi is the questionable measurement, 𝑋̅ is the mean, s is the standard deviation of the whole data set. The
suspected outlier can be rejected if g exceeds the tabulated g value. A table for critical g values for 3 to 10
measurements at 95% and 99% confidence level is given below:
MATERIALS
(10) 25-centavo coins Forceps or crucible tongs
GLASSWARE
Watch glass
EQUIPMENT
Analytical balance
PROCEDURE
Weighing of Samples
1. Place the ten coins on a watch glass using forceps.
2. Take the weight of each coin using "weighing by difference" method (refer to Appendix 2).
3. Record the weights of the coins in your data sheet.
EXPERIMENT 2
Solution Preparation
OBJECTIVES
At the end of the experiment, the student should be able to:
1) perform stoichiometric calculations needed in the solution preparation;
2) know the proper way of preparing solutions from solid and liquid reagents;
3) know the proper pieces of glassware and equipment to in solution preparation; and
4) calculate the exact concentration of the prepared solution from standardization.
INTRODUCTION
Many of the reactions utilized for quantitative analysis take place in aqueous solutions. In aqueous solution,
the species present are in a small state of subdivision, such that solutes exist as ions rather than aggregates of
these. In this manner, the particles are free to move about the solution and a proportion of the ensuing collisions
between solute particles result in a reaction. Such solutions may be prepared in two ways, depending on the
nature of the solute, whether solid or liquid.
The most common way of preparing solutions of different concentrations is by dissolving a weighed amount
of solid in enough solvent to produce a solution of specific volume. Weighing may be done with the use of a
top loading balance or an analytical balance depending on the accuracy sought for.
Another way of solution preparation is by dilution. This method requires measurement of a specific volume of
solute and rendering it less concentrated by addition of water until a solution of desired volume is achieved.
The most appropriate way of volume measurement is with the use of a graduated cylinder or a pipette
(volumetric/transfer or measuring), again, depending on the accuracy sought for. A volumetric pipette is used
for accurate measurements since it is designed and calibrated to deliver only one volume while a measuring
pipette is calibrated into small divisions allowing measurement of various amounts of liquid. Volumetric flasks
are also used to prepare the above mentioned solutions. The concentrations of such solutions may be expressed
in any manner appropriate, as discussed in the ways of expressing concentration, usually in molarity (M) or
parts per million (ppm).
In the laboratory, there are cases when the intended concentrations of the solutions to be used are too dilute
such that measurement of small masses or volumes seems impossible with the balances and glassware available.
In such situations, it is wiser for the analytical chemist to prepare a solution of relatively higher concentration
(a stock solution), measure out an appropriate volume (an aliquot) andthen further dilute it to prepare the
solution of desired (lower) concentration.
Weigh the desired amount in a beaker. Add enough distilled water to dissolve the solid then transfer
quantitatively to 100.0-mL volumetric flask. Dilute to mark with distilled water to yield 0.1000 M NaCl solution.
M1V1 = M2V2
(V1)(0.1000 M ) = ( 100.0 mL )( 0.0100 M )
V1 = 10.0 mL of the 0.1000 M NaCl stock solution prepared
measure out an aliquot of the stock solution and dilute to 100.0 mL (the desired volume) with water in a
volumetric flask.
For sample analysis, there are cases when the analyte present in the sample exists in high concentrations such
that dilution, as described above, is also necessary.
Twenty (20.00) mL of drinking water was measured and diluted to 100.0 mL. The diluted aliquot was analyzed via a
known titrimetric method (Experiment 7: Complexometric Titration). By applying the appropriate volumetric
calculation, the concentration of the analyte in the diluted sample (100.0 mL solution) may be determined.
NOTE: Expect that the calculated concentration of the original solution is HIGHER, (more concentrated). Notice how
the calculation resembles dimensional analysis. There are also cases when a series of dilutions have been performed such
that two or more DFs are necessary.
There are also cases when the concentration of the original solution is known and the concentration of a diluted solution
is to be determined (opposite of the example above). For this type of calculation, an ALIQUOT FACTOR may be utilized.
𝑣𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑎𝑙𝑖𝑞𝑢𝑜𝑡 1
𝑨𝒍𝒊𝒒𝒖𝒐𝒕 𝑭𝒂𝒄𝒕𝒐𝒓, 𝑨𝑭 = = (2.2)
𝑡𝑜𝑡𝑎𝑙 𝑣𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 𝐷𝐹
𝑌 𝑚𝑜𝑙𝑒𝑠 𝐶𝑎2+ 20.00 𝑚𝐿 𝑜𝑓 𝑠𝑎𝑚𝑝𝑙𝑒 𝑋 𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝐶𝑎2+
( )( )=
𝐿 𝑜𝑓 𝑠𝑎𝑚𝑝𝑙𝑒 0.10 𝐿 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 𝐿 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛
NOTE: Expect that the calculated concentration of the original solution is LOWER (less concentrated). Notice how the
calculation resembles dimensional analysis. There are also cases when a series of dilutions have been performed such
that two or more AFs are necessary.
The exact concentration of the prepared solution can then be determined by standardization. In standardizing
solutions, a primary standard is weighed with high accuracy, dissolved and then titrated with the solution until
the endpoint is reached. A primary standard should be a stable solid with high purity and high molecular
weight with a known chemical reaction with the solution to be standardized.
In this experiment, solutions of NaOH and HCl will be prepared to illustrate ways of solution preparation
described above. Afterwards, standardization of HCl and NaOH, using Na2CO3 and KHP as primary standards,
respectively, will then be done to determine the exact concentration of the standard solutions.
CHEMICALS
concentrated HCl NaOH KHP, primary standard
Na2CO3, primary standard Phenolphthalein
GLASSWARE
Volumetric flasks (100-, 250-mL) Beakers (250-mL) Droppers
Erlenmeyer flasks (250-mL) Volumetric pipette (25-mL) Spatula
EQUIPMENT
Analytical balance Top loading balance
PROCEDURE
NOTE: An analytical balance is not suitable in weighing NaOH pellets since it is hygroscopic.
3. Add enough distilled water to dissolve the pellets and stir. The dissolution is exothermic, so cool the
solution in a water bath if necessary.
4. When completely dissolved, transfer quantitatively into a 100-mL volumetric flask using distilled water to
wash the beaker. Add enough distilled water to make a volume of about 90.0 mL. Cover and cool the flask
and solution to room temperature.
5. Bulk the solution to the mark with distilled water and cover. Mix the solution thoroughly by repeated
shaking and inversion of the flask.
6. Transfer the solution into a dry and clean plastic bottle and label properly.
NOTE: Never store any solution in a volumetric flask as it is not a storage container. Never store NaOH
or any basic solutions in glass containers.
NOTE: Always add concentrated acid to water; never water to acid when diluting acid solutions. Add
enough distilled water to make a volume of about 40.0 mL. Swirl to mix, cover the flask and cool the
solution to room temperature, if necessary.
3. Bulk the solution to the mark with distilled water and cover. Mix the solution thoroughly by repeated
shaking and inversion of the flask.
4. Transfer the solution into a dry and clean reagent bottle and label properly.
EXPERIMENT 3
Iodine Clock Reaction
OBJECTIVES
At the end of the experiment, the student should be able to:
1) describe the kinetics of the reaction of I- and S2O82-;
2) use the initial rate method to determine the rate law of the reaction;
3) observe the effect of temperature on the reaction rate and calculate pertinent values of the Arrhenius
equation; and
4) observe the effect of a catalyst to the reaction rate.
INTRODUCTION
The rate of a chemical reaction is defined as the change of the concentration of a reactant or product per unit
time. The concentration of the reactants, temperature, and the presence of catalyst are the major factors that
affect the rate of a chemical reaction.
The kinetics of the reaction between persulfate, S 2O82-, and iodide, I- ions will be studied in this experiment:
The rate expression of this reaction can be written as the decrease of concentration of S 2O82-/ I- or the formation
of the products SO42-/ I2 with time. This is given in the following rate law:
where k is the rate constant and the powers x and y give the order of the reaction with respect to S2O82- and I-,
respectively. These variables can be determined experimentally using the method of initial rates. Such method
involves performing the reaction at controlled conditions, i.e. varying concentrations of one reactant while
keeping the concentration of the other constant, and measuring the rate at each case.
The effect of temperature on the reaction rate is given by the Arrhenius equation:
𝐸𝑎
ln 𝑘 = ln 𝐴 − (3.3)
𝑅𝑇
where A is the Arrhenius constant, Ea is the activation energy of the reaction, T is temperature in Kelvin, and R
is the universal gas constant (8.314 J mol -1 K-1). From the rate constants and reaction temperatures, Ea and A of
the reaction of S2O82- and I- can be determined.
In this experiment, the rate of I2 formation will be measured to describe the rate of the reaction. The I 2 formed
from the S2O82-/I- reaction is reduced back to I- by S2O32- ions.
When all the S2O32- is used up, free I2 starts to form in solution. By measuring the time taken for the known
amount of S2O32- to be consumed, the rate of the formation of I2 during that time can be calculated.
MATERIALS/APPARATUS
Ice Stopwatch/timer Thermometer
CHEMICALS
KI KCl K2SO4
K2S2O8 Na2S2O3∙5H2O CuSO4·5H2O
Starch
GLASSWARE
Volumetric flasks (5-, 500-mL) Measuring pipettes (10-mL) Beakers (50-mL)
EQUIPMENT
Top loading balance Hotplate
PROCEDURE
Solution Preparation
Prepare the following solutions by class:
1. 500.0 mL 0.2 M KI
2. 500.0 mL 0.2 M KCl
3. 500.0 mL 0.1 M K2S2O8
4. 500.0 mL 0.1 M K2SO4
5. 500.0 mL 4.0 mM Na2S2O3 (from Na2S2O3∙5H2O)
6. 20.0 mL of 1% (w/v) fresh starch solution
a. Moisten 0.20 g of soluble starch with a small amount of H2O until a smooth paste is obtained.
b. Pour slowly into 20.0 mL of boiling water. The starch solution must be freshly prepared.
NOTE: Keep the solution at 90°C-100°C to avoid starch solution from drying up.
7. 50.0 mL 0.01 M CuSO4 (from CuSO4·5H2O)
Table 3.1. The different runs for the effect of persulfate and iodide concentrations on reaction rate.
Beaker A Beaker B (+ 3 drops of fresh starch)
Run a
0.2 M KI, mL 0.2 M KCl, mL 0.1 M K2S2O8 0.1 M K2SO4 4.0 mM Na2S2O3
1 10.0 0.0 5.0 5.0 5.0
2b 5.0 5.0 5.0 5.0 5.0
3 2.5 7.5 5.0 5.0 5.0
4 5.0 5.0 7.5 2.5 5.0
5 5.0 5.0 10.0 0.0 5.0
a Condition: room temperature
b will be referred to as set 1 of run 2 (two more sets will be prepared in Part C)
2. Pour the contents of beaker A into beaker B. Immediately start timing the reaction. Stop the timer once the
mixture turns blue. Measure the temperature of the reaction mixture.
NOTE: Runs 1-5 SHOULD be done one at a time. For each run, label two 50-mL beakers as “A” and “B.”
WASTE DISPOSAL
1. Dispose of Cu(II) solutions into the inorganic waste jar.
2. All other solutions can be discarded in the sink along with copious amounts of water.
EXPERIMENT 4
Common Ion Effect and Buffers
OBJECTIVES
At the end of the experiment, the student should be able to:
1) understand and relate the concepts of common-ion effect and buffer solutions;
2) distinguish buffer solutions from other types of solutions;
3) perform calculations related to the buffer concept; and
4) observe the effect of dilution on the pH of a buffered sample.
INTRODUCTION
Acid-base indicators are weak acids or bases whose conjugate ions have different colors with respect to their
neutral molecules. Consider a hypothetical indicator, HInd, ionizing as follows,
where HInd is a weak acid with Color A and Ind- is the conjugate base with Color B. Within a certain pH range,
approximately equal amounts of HInd and Ind - are present. Thus, the solution will have a color intermediate
between Color A and Color B. Below this pH range, the HInd form predominates and the solution will exhibit
Color A. Above this range, the form Ind- predominates and the solution will show Color B. For instance,
indicators such as methyl orange and phenolphthalein have different colors at some pH ranges.
The pH of a solution can be estimated by adding a few drops of an indicator solution to it and observing the
resulting color. The exact pH may be determined by means of a calibrated pH meter (refer to Appendix 2).
In accordance with the Le Chatelier’s principle, addition of the products of a reaction to a system at equilibrium
causes the equilibrium to shift towards the formation of the reactants. Consider HA as a weak acid and thus
the equilibrium,
is shifted in the direction of HA by the addition of a common ion. Addition of a strong acid increases H 3O+ or
addition of a salt containing anion A - suppresses the ionization of HA.
Similarly, the addition of a strong base or a salt containing the cation BH + to a solution of B, a weak base shifts
the equilibrium towards B, lowering the degree of ionization of B
A buffer is a solution which resists an appreciable change in pH upon addition of small amounts of strong acid
or strong base. An example of this is a solution which contains a weak acid or weak base with their respective
conjugate ions. Given this condition, a solution containing B and BH + is a buffer. Likewise, a solution containing
HA and A- is also an example. Given below are the equations which show the buffer action of:
1. HA–A - Buffer
To illustrate how a buffer works, consider a buffer composed of a weak acid (HX) and one of its salts (MX,
where M+ could be Na+, K+, or other cations). The acid-dissociation equilibrium in this buffered solution involves
both the acid and its conjugate base:
[𝐻3 𝑂 + ][𝑋 − ]
𝐾𝑎 = (4.11)
[𝐻𝑋]
𝐾𝑎 [𝐻𝑋]
[𝐻3 𝑂 + ] = (4.12)
[𝑋 − ]
It is seen from this expression that [H3O+], and thus the pH, is determined by two factors: the value of Ka for the
weak-acid component of the buffer, and the ratio of the concentrations of the conjugate acid-base pair, [HX] /
[X-]
If OH- ions are added to the buffered solution, they react with the acid component of the buffer
This reaction causes [HX] to decrease and [X-] to increase. However, as long as the amounts of HX and X- in the
buffer are large compared to the amount of OH- added, the ratio [HX] / [X-] doesn’t change much, and thus the
change in pH is small.
If H3O+ ions are added, they react with the base component of the buffer:
Using either of the two equations, it is observed that the reaction causes [X -] to decrease and [HX] to increase.
As long as the change in the ratio [HX] / [X-] is small, the change in pH will be small.
Buffer capacity is the amount of acid or base the buffer can neutralize before the pH begins to change to an
appreciable degree. The capacity of a buffer is within a certain range: pH = pKa ± 1 or pOH = pKb ± 1. This
capacity depends on the amount of acid and base from which the buffer is made.
The pH of a buffer depends on the Ka for the acid and on the relative concentrations of the acid and base that
comprise the buffer. An alternate approach used to calculate the pH of a buffer is the Henderson-Hasselbalch
equation:
[𝑏𝑎𝑠𝑒]
𝑝𝐻 = 𝑝𝐾𝑎 + 𝑙𝑜𝑔 ( ) (4.15)
[𝑎𝑐𝑖𝑑]
where [acid] and [base] refer to the initial concentrations of the acidic and basic components of the buffer,
respectively. Henderson-Hasselbalch equation assumes that the initial concentrations of the acidic and basic
components are approximately equal to their equilibrium concentrations. That is, the dissociation of the acidic
component is negligible.
[𝑎𝑐𝑖𝑑]
𝑝𝑂𝐻 = 𝑝𝐾𝑏 + 𝑙𝑜𝑔 ( ) (4.16)
[𝑏𝑎𝑠𝑒]
CHEMICALS
CH3COOH NH3 HCl
NaOH NaCH3COO NH4Cl
Phenolphthalein Methyl orange
GLASSWARE
Volumetric flasks (250-,100-mL) Beakers (50-mL) Measuring pipettes (10-mL)
Volumetric pipettes
EQUIPMENT
pH meter
PROCEDURE
Solution Preparation
Prepare the following solutions by group:
1. 50.0 mL 0.20 M CH3COOH
2. 50.0 mL 0.10 M CH3COOH
3. 50.0 mL 0.20 M NH3
4. 50.0 mL 0.10 M NH3
5. 25.0 mL 0.20 M NaCH3COO
6. 25.0 mL 0.20 M NH4Cl
WASTE DISPOSAL
1. Discard excess concentrated acids and bases into their respective waste bottles.
2. Drain and flush excess diluted solutions into the sink with copious amounts of water.
EXPERIMENT 5
Determination of the Solubility Product Constant of
Calcium Hydroxide
OBJECTIVES
At the end of the experiment, the student should be able to:
1) determine the Ksp of calcium hydroxide; and
2) explain the effects of common and diverse ions on solubility of sparingly soluble salts.
INTRODUCTION
When a slightly soluble ionic solid is placed in water, equilibrium is soon established between the excess solid
and the ions in the saturated solution. For example, the hypothetical ionic solid A xBy is in contact with its
saturated solution, the equilibrium is:
where the equilibrium constant, Ksp, is called the solubility product constant of AxBy(s) and the expression [Ay+]x
[Bx-]y is the ion-product, IP, or the reaction quotient, Q, when the concentrations used are initial concentrations.
Compared to other equilibrium expressions, it is noticeable that the solid itself does not appear as a
denominator in the Ksp expression. It is because the activity of any solid is equal to one. Thus, writing the K sp
expression for saturated and heterogenous solution, the concentration of solids is ignored. However, like all
equilibrium constants, Ksp is temperature dependent.
The Ksp value may be used as a measure of solubility, and the relation between IP of the solution and the K sp
can be used as a criterion if precipitation will happen or not. When A y+ ions were added to a solution containing
Bx- ions:
Other factors, such as presence of common ions in the solution and ionic strength of the medium, also affect the
molar solubility of a compound. A common ion is an identical ion from another substance that is also present
in the solution. According to Le Chatelier’s principle, the presence of a common ion shifts the direction of the
equilibrium to the reactant side. Therefore, the concentration of the common ion should be taken into account
when determining solubility.
However, when other ion or electrolyte from a substance containing no common ion with the sample is present
in the solution, the resultant increase in the ionic strength of the solution is what is accounted in the solubility
of the compound.
The ionic strength, μ, of the solution depends on the charges and the concentrations of all the ions present in
the solution,
1
𝜇= ∑ 𝐶𝑖 𝑍𝑖 2 (5.3)
2
where Ci and Zi are concentrations (in molarity) and charges of the ions, respectively.
1
μ= (C + (Z𝑁𝑎+ )2 + C𝑆𝑂42− (Z𝑆𝑂42− )2 )
2 𝑁𝑎
1
μ = ((0.20)(1)2 + (0.10)(−2)2 )
2
μ = 0.30
In the experiment, the Ksp of calcium hydroxide, Ca(OH)2 will be determined by calculating for the [OH-] from
a saturated solution of Ca(OH)2 through acid-base titration.
The effect of common ion and change in ionic strength of the solution to the solubility of Ca(OH) 2(s) will also
be evaluated.
APPARATUS
Iron stands & rings Burette clamps Filter paper
CHEMICALS
HCl, concentrated Ca(OH)2 Phenolphthalein
Na2CO3, primary standard KCl CaCl2
GLASSWARE
Beakers (100-, 250-mL) Measuring pipettes Volumetric pipettes (10-, 25- mL)
Erlenmeyer flasks (250-mL) Burettes (50-mL) Volumetric flasks (250-,1000-mL)
Graduated cylinder (100-mL) Funnels
EQUIPMENT
Analytical balance Top loading balance Hotplate
PROCEDURE
Solution Preparation
Prepare the following solutions by class:
1. 1.00 L 0.10 M HCl
NOTE: Use the 3.0 M HCl solution from the Solution Preparation experiment.
2. 250.0 mL 1.0 M stock KCl
3. 250.0 mL each 0.50 M, 0.25 M, 0.10 M, 0.05 M, 0.01 M KCl solution from stock 1.0 M KCl solution.
4. 250.0 mL 0.010 M CaCl2
WASTE DISPOSAL
1. Collect excess HCl solution. Use this to dissolve all Ca(OH)2 precipitate. Dilute the resulting solution with
plenty of water and flush directly down the sink with copious amounts of water.
2. Dispose of all titrated solution into the sink with copious amounts of water.
3. Dispose of all used filter papers in the solid waste container.
EXPERIMENT 6
Quantitative Determination of Soda Ash Composition by
Double Indicator Titration
OBJECTIVES
At the end of the experiment, the student should be able to:
1) determine the composition of the soda ash sample and respective percentage in the sample,
2) relate the experiment to the following concepts: (a) strong acid-strong base and strong acid-weak base
titrations, (b) carbonate system titration, and (c) double-indicator titration, and
3) perform calculations involving simple and complex acid-base titrations and in particular those dealing
with carbonate-like systems.
INTRODUCTION
Volumetric or titrimetric methods are processes that involve the reaction of a standard solution of known
amounts with an unknown solution or analyte to determine the stoichiometric or equivalence point. If the
details of the reaction are well-defined and the equivalence point is accurately located, then the amount of the
unknown or analyte present can be calculated from the known amount of the standard solution used in the
reaction.
In any volumetric analysis, a reference material is most important as the accuracy of the method is dependent
on it. This reference material is the primary standard upon which one determines the accurate concentration
of the standard solution.
Standard titrants in acid-base titrations are generally strong acids or strong bases. In aqueous solutions, weak
acids and bases are not suitable due to the pH change that takes place near the equivalence point. The most
commonly used acid is hydrochloric acid and for the base, sodium hydroxide.
Acid-base titration is based on neutralization reaction, this is also called acidimetric or alkalimetric titration. It
is a simple reaction of a proton and a hydroxyl ion as given by the ionic reaction:
Simple systems usually involve monofunctional acids or bases. Mixtures of acids or bases and polyfunctional
acids or bases are considered complex systems. Such complex systems usually contain two or more acidic or
basic species.
In this experiment, titration of complex systems will be observed. With hydrochloric acid as titrant, the primary
standard used for the standardization is sodium carbonate with high purity. The carbonate ion reacts with the
acid according to the successive acid-base reactions:
It is a two component system wherein the second component is a product of the first reaction. The first
component is the carbonate ion (CO32-) and the second is the bicarbonate ion (HCO3-). Following stoichiometry,
it can also be deduced that the amount of H3O+ in reactions (6.2) and (6.3) are equal. The first equivalence point
(6.2) can be seen using phenolphthalein indicator with a color transition from pink to colorless. Moreover, the
second equivalence point (6.3) can be seen using methyl orange indicator.
In this analysis, the unknown sample may be a pure compound of sodium carbonate, sodium bicarbonate or
sodium hydroxide or a compatible mixture of the three. A mixture of sodium hydroxide and sodium
bicarbonate are incompatible due to their properties. The analysis of the unknown sample will need a two-
titration or a double indicator titration. In this case, a double indicator titration will be used. The indicators to
be used should change color at different pH ranges, to signal neutralization of different protons of the complex
system. The first reaction involves reactions (6.1) and (6.2) where the phenolphthalein indicator changes color;
while the second reaction involves reaction (6.3) where the methyl orange indicator changes color. Thus, to
resolve the possible components of an unknown sample requires comparison of the volume of HCl required
reaching these two distinct endpoints, the Vph and Vmo, as shown in Figure 6.1.
Figure 6.1. A graph of pH against volume of HCl titrant for reactions involving carbonate and carbonate-like
systems.
Table 6.1 gives the volume relationships in the analysis of soda ash using double indicator titration method.
Table 6.1. The volume relationships for different compositions of soda ash.
Substance composition Relation between Vph and Vmo Amount of substance present
NaOH Vmo = 0; Vph > 0 MVph
Na2CO3 Vph = Vmo MVph or MVmo
NaHCO3 Vph = 0; Vmo > 0 MVmo
NaOH + Na2CO3 Vph > Vmo NaOH: M(Vph – Vmo)
Na2CO3: MVmo
NaHCO3 + Na2CO3 Vph < Vmo NaHCO3: M(Vmo – Vph)
Na2CO3: MVph
In this experiment, a sample of soda ash will be analyzed using the double indicator titration described.
APPARATUS
Iron stands Burette clamps
CHEMICALS
HCl, concentrated Na2CO3, primary standard Soda ash sample
Methyl orange Phenolphthalein
GLASSWARE
Burettes (50-mL) Volumetric flasks (250-mL) Measuring pipettes (25-mL)
Volumetric pipettes (25-mL) Erlenmeyer flasks (250-mL) Beakers
Graduated cylinders (100-mL) Spatulas
EQUIPMENT
Analytical balance Hotplate
PROCEDURE
NOTE: The use of boiled distilled water is necessary for this experiment. Boil 2.00-L distilled water a day before
the experiment. Store this in a sealed container.
NOTE: Record all data with tolerances. Account these for error propagation. For error propagation instructions,
refer to Appendix 6.
Preparation of Solutions
Prepare the following solution quantitatively using boiled distilled water per class:
1. 250.0 mL 1.0 M stock HCl solution
NOTE: Use the 3.0 M HCl solution from Solution Preparation experiment
Prepare the following solution quantitatively using boiled distilled water per group:
1. 250.0 mL 0.0500 M standard HCl solution from 1.0 M HCl
3. Resume titration to the methyl orange endpoint, which is the formation of an orange solution.
4. Record the final volume of acid at the methyl orange endpoint.
WASTE DISPOSAL
1. Dispose of all titrated solutions into the sink with copious amounts of water.
2. Dispose of excess stock 1.0 M HCl into the acid waste container.
CALCULATIONS:
Determine the:
1. molarity of the standard HCl solution and report it as M ± ∆M
2. percentage composition of the sample and report it as %A ± ∆%A
3. relative standard deviation (in ppt) and confidence limits (95% confidence level)
EXPERIMENT 7
Quantitative Determination of Total Hardness in Drinking
Water by Complexometric EDTA Titration
OBJECTIVES
At the end of the experiment, the student should be able to:
1) apply the concept of complexometric titration in the determination of total hardness in drinking water.
INTRODUCTION
One of the most common analyses for water, whether it is for domestic (potable) or industrial (process) use, is
total hardness. This parameter is basically associated to the amount of calcium and magnesium ions in water.
Monitoring of this is important because high amount of calcium and magnesium ions in water will cause several
problems. First, these ions form precipitates with soaps which will lessen the cleansing action of the soap. Also,
hard water would precipitate calcium carbonate and magnesium carbonate on boiling resulting to clogging of
pipes of the boiler equipment.
The total hardness of water samples is usually determined by complexometric EDTA titration using Eriochrome
Black T (EBT) as indicator and is reported as ppm calcium carbonate. Using the result in ppm CaCO 3, one can
classify the hardness of water based on the table below.
In this experiment, the amount of calcium and magnesium in water samples will be determined by titration
with ethylenediaminetetraacetic acid (EDTA). The endpoint takes place when EDTA reacts with the colored
metal-indicator complex, thus breaking the complex. The titration is maintained at pH 10 to allow the Ca-EDTA
and Mg-EDTA complex to form stoichiometrically. The wine-red color of the MgIn- complex breaks up at the
equivalence point as illustrated by the reaction:
APPARATUS
Iron stands Burette clamps
CHEMICALS
CaCO3, primary standard Na2H2EDTA●2H2O MgCl2.6H2O
HCl, concentrated NH3, concentrated NaOH
NH4Cl EBT (1:100:100 / w:w:w /EBT:NaCl:NH2OH.HCl)
GLASSWARE
Volumetric flasks (50-, 100-, 250-, 500-mL) Measuring pipettes (5-mL)
Volumetric pipettes (5-, 10-, 25-, 50-mL) Beakers (250, 400-mL)
Erlenmeyer flasks (250-mL) Burettes (50-mL) Watch glass
EQUIPMENT
Analytical balance Top loading balance Hotplate
pH meter
PROCEDURE
Preparation of Solutions
Prepare the following solutions by class:
1. 500.0 mL 0.1000 M stock EDTA solution
a. Weigh an appropriate amount of Na2H2EDTA2H2O (FW=372.24) to the nearest 0.1 mg and transfer
to a 400-mL beaker.
b. Add about 200 mL distilled water. Stir to dissolve then add 1.0 g MgCl 26H2O crystals. Mix to
dissolve the crystals.
NOTE: The dissolution of Na2H2EDTA∙2H2O may be slow so addition of NaOH pellets while stirring
may be added until the solution is clear. Heating the solution may also increase the dissolution of
Na2H2EDTA∙2H2O.
c. Quantitatively transfer the solution into a 500-mL volumetric flask. Rinse the beaker thrice with
small portions of distilled water and transfer the rinse to the volumetric flask. Dilute to mark with
distilled water, cover and mix.
d. Transfer to a clean and dry reagent bottle. Label as 0.1000 M stock EDTA solution.
NOTE: The stirring must be vigorous enough to dissolve the solids. It is advisable to swirl the beaker
using a crucible tong instead of using a stirring rod.
c. Quantitatively transfer the solution into a 100-mL volumetric flask. Dilute to mark with distilled
water, cover and mix.
d. Transfer to a clean, plastic (polyethylene) bottle. Label as 0.0500 M stock Ca2+ solution.
NOTE: Avoid adding too much indicator, otherwise the endpoint will not be sharp, with a gradual color
change that is difficult to detect.
3. Record the initial volume of the titrant. Titrate the solution until a change in color from wine red to a clear
blue is observed.
NOTE: Titrate the solution slowly as the endpoint approaches because the color change is delayed. Over
titration may result from rapid titration.
NOTE: The color change may also be observed as from light violet/purple to blue.
WASTE DISPOSAL
1. Dispose of all titrated solutions into the sink with copious amounts of water.
2. Drain and flush excess solutions into sink with copious amount of water.
EXPERIMENT 8
Quantitative Determination of Dissolved Oxygen Content
by Winkler Redox Titration
OBJECTIVES
At the end of the experiment, the student should be able to:
1) perform the water sampling and pre-treatment techniques for dissolved oxygen analysis;
2) determine the amount of dissolved oxygen in a water sample from a pond in the University using
Winkler redox titration; and
3) discuss the chemistry behind the Winkler method for dissolved oxygen determination.
INTRODUCTION
The dissolved oxygen (DO) levels in natural water and wastewater depend on the physical, chemical and
biochemical processes involved in the water system. Dissolved oxygen determination is a key test for water
pollution control and waste water treatment process control. Table 8.1 shows the water quality guidelines, as
per American Public Health Association, Inc.
Winkler method, the classical determination of DO in water is based on an oxidation-reduction titration process
known as iodometric method. The basis of this method is the oxidizing power or ability of the dissolved oxygen
to oxidize the divalent manganese in the solution. The oxidized Mn is precipitated to hydroxides of higher
valence states (Mn2+ → Mn3+) as Mn(OH)3 with strong alkali. Upon acidification, the oxidized Mn(III) is reduced
to Mn(II) in the presence of iodide ions with subsequent liberation of iodine equivalent to the DO content of the
water sample.
Several oxidizing and reducing substances, such as dissolved organic matter, nitrate, nitrite, higher valence
manganese compounds, active chlorine, sulfide, sulfite, ferrous, and ferric, may be present in natural water or
wastewater. Due to the possible presence of these substances, water sampling has been intricately studied and
an accepted technique should be followed before performing DO analysis. Proper sampling techniques, which
include pre-treatment procedures, are performed to avoid interferences of these substances.
APPARATUS/MATERIALS
Iron stands Burette clamps Aluminum foil
CHEMICALS
MnSO4·2H2O NaOH NaN3
KI Na2S2O3·5H2O Na2CO3
H3PO4 KIO3, primary standard
GLASSWARE
Glass bottle with cap Watch glass volumetric flasks (25-, 250-, 500- mL)
Burettes (50-mL) Syringes (1-mL) Na2CO3
Beakers (50-, 100-mL) Erlenmeyer flasks (250-mL)
EQUIPMENT
Analytical balance Top loading balance Hotplate
PROCEDURE
Solution Preparation
Prepare the following solutions by class:
1. 25.0 mL 4.0 M MnSO4 from MnSO4·2H2O crystals
a. Weigh the needed amount of crystals in a beaker and dissolve in about 10 mL distilled water.
b. Filter the solution into a 25-mL volumetric flask and dilute to mark.
Standardization of Na2S2O3
1. In a 50-mL beaker, weigh 0.15 g of the primary standard KIO3 to the nearest 0.1 mg. Record the weights
of the primary standard.
2. Dissolve in about 50 mL of distilled water.
3. Transfer the solution quantitatively into a 100-mL volumetric flask. Dilute to mark and mix thoroughly.
4. Take three 10.00 mL aliquots and transfer into three 250-mL Erlenmeyer flasks.
5. Add about 20 mL distilled water into each flask.
6. Add 1.0 g KI and 10 mL of 0.5 M H2SO4 to each solution.
7. Immediately titrate the solution with standard Na2S2O3 until a pale yellow color is obtained.
8. Immediately add 1.0 mL starch solution.
9. Continue titration until the disappearance of the blue color.
10. Perform standardization in triplicate.
3. Close the bottle carefully, avoiding inclusion of air bubbles. Shake the bottle thoroughly and vigorously.
At this point, the solution is filled with precipitates.
4. Remove the cover slowly and add 2.0 mL of concentrated H 3PO4 taking care that the pipette must be just
below the surface of the water. The acid will dissolve the precipitate.
5. Cover and shake the solution and allow it to stand for about 10 minutes.
6. Take a 50.0 mL aliquot of the solution and transfer in a 250 mL Erlenmeyer flask.
7. Follow steps 7-9 of standardization.
8. Perform sample analysis in triplicate.
9. Express the DO content in the samples in parts per million (ppm).
WASTE DISPOSAL
1. Dispose of all titrated solutions into the sink with copious amounts of water.
2. Excess H2SO4, Na2S2O3, and starch solutions may be discarded into the sink with copious running water.
3. Dispose of excess NaOH with KI and NaN3 into the base waste container.
4. Dispose of excess MnSO4 into the inorganic waste container.
EXPERIMENT 9
Determination of Electrode Potentials
OBJECTIVES
At the end of the experiment, the student should be able to:
1) relate and apply the concepts of electrochemistry to actual experiments;
2) understand the processes and elements of an electrochemical cell; and
3) determine the spontaneity of redox reactions based on standard reduction potential.
INTRODUCTION
Reactions in which one or more electrons are transferred are called reduction-oxidation reactions or redox
reactions. Reduction involves gain of electrons while oxidation involved loss of electrons. For instance, direct
transfer of electrons from Zn atom to each Cu2+ ion occurs when a piece of Zn metal is dropped into a CuSO4
solution as shown by the equation (9.1).
The reaction is spontaneous but no useful electrical work is performed. However, if the two reactants are
separated in such a way that electron transfer is forced through a wire (Figure 9.1); electrical work can be done
by the reaction system. Such a device is called galvanic or voltaic cell.
Reaction 9.1 can be split into two half-reactions (9.2 and 9.3) representing two half-cells:
In Figure 9.1, the two half-cells are connected by a wire with a voltmeter and a salt bridge. Electrical current
flows from the Zn electrode (anode) to the Cu electrode (cathode) through the voltmeter and the circuit is
completed by migration of ions thru the salt bridge: Na+ toward the cathode and SO42- toward the anode. Thus,
electroneutrality of the two solutions is maintained.
The Cell Potential, Ecell, and the Standard Cell Potential, E0cell
To conveniently describe a galvanic cell, a shorthand notation, called the cell notation, is used instead of
drawing the complete diagram. In a cell notation, the components of the cell are written according to the
movement of electrons: from anode to cathode. A single bar indicates a phase boundary while the double bar
represents the salt bridge. Figure 9.1 can be represented by the cell notation:
Such cell approximates what is known as standard cell where all substances are at unit activity. In a galvanic
cell, the oxidizing agent pulls electrons from the reducing agent through the conduction wire. This electron pull
or driving force is called the cell potential (Ecell) or the electromotive force, emf of the cell. However, at standard
conditions, the potential is called the standard cell potential (E0cell). Ecell is a measure of the tendency for a cell
reaction to occur for the conditions under which the cell operates while E0cell measures the tendency for the cell
reaction to occur when all substances involved are at 1.0 M concentration.
The dependence of cell potential on concentration results directly from the dependence of free energy on the
concentration which leads to the Nernst equation (9.4 and 9.5). The Nernst equation gives the relation of the cell
potential, Ecell, to the nature of the electrodes, temperature, and concentration of substances involved in the cell
reaction.
2.303𝑅𝑇
𝐸𝑐𝑒𝑙𝑙 = 𝐸 𝑜 𝑐𝑒𝑙𝑙 − log 𝑄 (9.4)
𝑛𝐹
𝑅𝑇
𝐸𝑐𝑒𝑙𝑙 = 𝐸 𝑜 𝑐𝑒𝑙𝑙 − ln 𝑄 (9.5)
𝑛𝐹
where R=8.314 J/mol-K; T=temperature in K; n=moles of electrons transferred; F=96,485 C/mol electrons
transferred, and Q= reaction quotient. Solids and pure solvents are assigned unit activity.
is set up. The tendency of a metal ion to form the metal atom (9.6) differs from one metal to another. Thus, the
potential set up between the metal and its ions varies from one system to another. The potential of a half-cell is
a measure of the tendency of the half-cell reaction to occur. The greater the reduction potential is (i.e. the more
positive), the greater the tendency for the reduction to occur and the lower the tendency for the reverse reaction
to occur. The more negative the reduction potential the lower the tendency for the reduction reaction to occur
and the greater the tendency for the reverse reaction (oxidation) to take place spontaneously. Thus, if two half-
cells are connected to form a cell, the half-cell with greater reduction potential will be the cathode and the one
with the lower reduction potential will be the anode. The cell potential is given by equations (9.7) and (9.8).
Standard reduction potentials, E0red, of half-cells are measured relative to the standard hydrogen half-cell which
has standard reduction potential of 0.00V arbitrarily assigned to it, i.e., for
The standard hydrogen half-cell can be replaced by half-cells whose standard reduction potentials have been
measured accurately against the standard hydrogen electrode.
In the first part of this experiment, standard reduction potentials of various half-cells will be measured against
the Cu2+(1M)/Cu half-cell whose standard reduction potential is given.
In the second part, electrolytic cell will be constructed. Electrolysis involves forcing a current through a cell to
produce a chemical change. This is different from the galvanic cell wherein the latter converts chemical energy
into electrical energy. Electrolysis will be conducted before measuring the standard reduction potentials of the
halide half-cells. To determine the amount of solid produced after the electrolysis, the Faraday’s Law (10.11)
may be used.
𝑄𝑀
𝑚= (9.11)
𝐹𝑧
where m is the mass of the substance liberated at the electrode in grams, Q is the total electric charge passed
through the substance, F is the Faraday’s constant equal to 96,485 C/mol e -, M is the molar mass of the substance,
and z is the number of electrons transferred per ion. Alternatively, the total electric charge can be determined
from the measured current (I) during the electrolysis multiplied by the duration (time, t) of the process in
seconds. This translates to equation (9.12).
𝐼𝑡𝑀
𝑚= (9.12)
𝐹𝑧
APPARATUS/MATERIALS
Copper wire Filter paper Alligator clips
Zinc stick Pencil graphite leads 9V or AA batteries
CHEMICALS
KI FeSO4·7H2O KBr
H2O2 FeCl3 ZnSO4·7H2O
KNO3 CuSO4·5H2O
GLASSWARE
Beakers (50-mL) Volumetric flasks (10-, 50-mL) Volumetric pipettes (5-mL)
Measuring pipettes (5-mL)
EQUIPMENT
Top loading balance Magnetic stirrer with spin bar Multimeter
PROCEDURE
Preparation of Solutions
Prepare the following solutions by group:
1. 50 mL 1.0 M CuSO4 from CuSO4·5H2O
2. 10 mL 1.0 M ZnSO4 from ZnSO4·7H2O
3. 10 mL 2.0 M FeCl3
4. 10 mL 2.0 M FeSO4 from FeSO4·7H2O
5. 10 mL 1.0 M Fe2+/Fe3+ solution
Pipette 5.0 mL each of 2.00M FeSO4 and 2.00 M FeCl3.
6. 10 mL 1.0 M KBr
7. 10 mL 1.0 M KI
8. 50 mL saturated KCl
NOTE: To check if the multimeter is working, set the reading to voltage. Then, connect the positive test
probe of the multimeter to the positive pole of the battery and the negative test probe of the multimeter to
the negative pole of the battery. The display should give a voltage reading. Otherwise, replace the
multimeter.
NOTE: Check if the alligator clips are working by inserting these in the circuit used in the previous step.
Connect the clips to the probes of the multimeter before connecting to the poles of the battery. Discard and
replace clips if the display does not give a voltage reading.
3. Fill one electrolyte container with 10 mL 1.0 M CuSO4 solution. Fill another electrolyte container with 10
mL 1.0 M ZnSO4. Prepare the set up as shown in Figure 9.2.
Cu electrode
(Cu wire) Zn electrode
Salt bridge (Zn stick)
Zn2+ solution
Cu2+ solution
4. Set the multimeter in voltmeter. Measure the voltage (Ecell). Note that the voltmeter should be PARALLEL
(Figure 9.2) to the circuit. The positive test probe of the multimeter is connected to the positive electrode,
while the negative test probe is connected to the negative electrode. A negative voltage reading will only
indicate a reversed attachment to the test probes, but the magnitude remains the same.
5. Repeat steps 3 and 4 for 1.0 M Fe2+/Fe3+, replacing the Zn2+ solution and use graphite as the electrode (Figure
9.3). The Cu2+/Cu half-cell and salt bridge must always be fresh per run.
Cu electrode
Graphite electrode
(Cu wire)
Salt bridge (Pencil lead)
Fe2+/Fe3+ solution
Cu2+ solution
Electrolytic Cells
1. Fill one electrolyte container with 10 mL 1.0 M KBr solution.
2. Prepare the set-up according to Figure 9.4.
3. Set the multimeter in ammeter. Connect the dry cell and ammeter in SERIES (Figure 9.4). Connect the
positive test probe of the ammeter nearer to the positive of the power supply, and the negative test probe
to the negative pole of the power supply. Simultaneously turn on the stirrer and start the timer upon
connecting the dry cell. Electrolyze or generate Br2 for 1 minute. Record any observations.
Graphite electrode
Spin bar
(Pencil leads)
Magnetic stirrer
4. Record the current during electrolysis. After the electrolysis, record the exact time elapsed. Be careful not
to disturb the container containing the Br2.
5. Set up another galvanic cell for the electrolyzed solution, as shown in Figure 9.5. Use fresh CuSO4 solution
and salt bridge. Measure the voltage (Ecell).
6. Do the same for 1.0 M KI solution.
Cu electrode
(Cu wire)
Salt bridge
WASTE DISPOSAL
1. Dispose of all potassium salt solutions into the sink with copious amounts of water.
2. Dispose of H2O2 in peroxide waste jar.
3. Dispose of all other used and excess reagents into the inorganic waste container.
EXPERIMENT 10
Quantitative Determination of the Purity and Dissociation
Constant of Potassium Hydrogen Phthalate by
Potentiometric Titration
OBJECTIVES
At the end of the experiment, the student should be able to:
1) discuss the principles involved in potentiometric titration;
2) detect the equivalence point in a titration curve using this method;
3) determine the purity of KHP; and
4) evaluate the acid dissociation constant of KHP from potentiometric data.
INTRODUCTION
Strong acids completely dissociate in water, but weak acids, like acetic acid, are only partially dissociated. For
a weak monoprotic acid represented by the formula, HA, partial ionization establishes the following
equilibrium:
HA ↔ H+ + A– (10.1)
[𝐴− ][𝐻 + ]
𝐾𝑎 = (10.2)
[𝐻𝐴]
Hence,
𝐾𝑎 [𝐻𝐴]
[𝐻 + ] = (10.3)
[𝐴− ]
In this experiment, the purity and acid dissociation constant of KHP, a weak acid will be determined using a
technique called potentiometric titration.
Potentiometric methods are analytical methods based upon potential measurements. Direct potentiometric
measurements compare the potential developed in a cell containing the indicator electrode in the analyte
solution with the potential developed when the indicator electrode is immersed in one or more standard
solutions of known analyte concentrations. The determination of pH using a glass membrane electrode is an
example of direct potentiometry.
V’ and V” are average volumes of the titrant used and are written in the x-axis of the first and second derivative
plots, respectively.
CHEMICALS
NaOH KHP, primary standard KHP sample
phenolphthalein
GLASSWARE
Burettes (50-mL) Erlenmeyer flasks (250-mL) Beakers (250-mL)
Graduated cylinders (100-mL)
EQUIPMENT
Analytical balance Top loading balance pH meter
Magnetic stirrer with spin bar
PROCEDURE
Solution Preparation
The following solution is prepared by group:
1. 250 mL 0.10 M NaOH
NOTE: Use the 1.0 M NaOH prepared from Solution Preparation experiment.
4. Switch on the magnetic stirrer and make sure that the spin bar does not hit the electrodes. Adjust the
electrode’s position if necessary.
5. Record the pH of the solution before starting the titration. Titrate the first sample solution by adding an
increment of 1.0 mL of the base. Record the volume of the base and the corresponding pH reading after
each addition of the titrant. Approximate the equivalence point by noting the volume at which a large
change in pH occurs.
NOTE: Use a spreadsheet software (e.g., Microsoft Excel) to record potentiometric titration data. Provide
the instructor with the file.
6. For the second and third trials, titrate the solution very carefully:
a. add the base solution using 1.0 mL increment at the beginning
b. at 5.0 mL of the equivalence point, titrate the sample solution using 0.5 mL increments
c. at 3.0 mL of the equivalence point, titrate the sample solution using 0.2 mL increments
d. at 2.0 mL of the equivalence point, titrate the sample solution using 0.1 mL increments.
7. Continue the titration beyond 5.0 mL of the equivalence point using 0.5 mL increments until pH 11 is
reached.
NOTE: Store the excess 0.10 M NaOH in reagent bottle for Ion Exchange Chromatography experiment.
WASTE DISPOSAL
1. Dispose of all titrated solutions into the sink with copious amount of water.
EXPERIMENT 11
Quantitative Determination of Copper(II) Concentration
by Spectrophotometry
OBJECTIVES
At the end of the experiment, the student should be able to:
1) apply spectrophotometry in the quantitative analysis of copper (II) solutions;
2) operate a spectrophotometer and measure transmission properties of solutions; and
3) determine an unknown copper (II) concentration in a sample using Beer’s Law.
INTRODUCTION
Spectrophotometry is an analytical technique in measuring transmission properties of materials as a function
of wavelength. The diagram below shows a radiant energy with intensity Io directed at a sample solution in a
transparent cell or cuvette. The solution can absorb a portion of this energy and the unabsorbed energy is then
transmitted with intensity I.
radiant energy /
light source
transmitted energy /
unabsorbed energy
Figure 11.1. Diagram of a beam of light as it travels through a sample in a cell of width l.
The amount of energy absorbed by the solution can be measured in transmittance, T, or in absorbance, A, where
𝐼
𝑇= (11.1)
𝐼𝑜
and
𝐼𝑜 1
𝐴 = log = log (11.2)
𝐼 𝑇
To determine the relationship of the absorbed energy to a solution’s concentration, one must apply the Beer’s
Law, or more accurately the Beer-Lambert-Bouguer Law,
𝐴 = 𝑎𝑏𝑐 (11.3)
where a = absorptivity with units of ppm-1. cm-1, a measure of how well a substance absorbs light
b = path length with units of cm, the width of the cell or cuvette in which the sample is contained
c = concentration of the component of interest in the solution with units of ppm
The Beer’s Law gives a convenient linear relationship of the absorbed energy with the concentration of the
absorbing species dispersed in the solution. In this experiment, Cu(II) solutions of known concentration are
prepared and are converted to [Cu(NH3)4]2+ species by the addition of concentrated ammonia. The absorbed
energies of these solutions are then collected at the pre-determined wavelength of maximum absorption, max.
Figure 11.2. Block diagram illustrating the components of a single-beam UV-Vis Spectrophotometer.
Furthermore, the various absorbance values of known Cu(II) concentration are graphed and correlated to
determine the unknown concentration of Cu(II) in a sample.
CHEMICALS
Cu(NO3)2.5H2O concentrated NH3
GLASSWARE
Volumetric flasks (50-, 250-mL) Measuring pipettes (10-mL) Beaker (100-mL)
EQUIPMENT
Top loading balance UV-Vis spectrophotometer 1 cm cuvette
PROCEDURE
Solution Preparation
1. 250.0 mL standard 2500 ppm Cu(II) stock solution
a. Weigh and dissolve appropriate amount of Cu(NO3)2·5H2O crystals in enough distilled water.
b. Transfer quantitatively into a 250-mL volumetric flask and dilute to mark.
2. From the plot of absorbance against wavelength, determine the analytical wavelength, max, for the
analysis.
WASTE DISPOSAL
1. Dispose of all excess standards and samples into the inorganic waste container.
EXPERIMENT 12
Quantitative Determination of Total Ion Concentration by
Ion-Exchange Chromatography
OBJECTIVES
At the end of the experiment, the student should be able to:
1. discuss the principles behind ion-exchange chromatography and its use as a technique for separation,
and
2. determine the total ion concentration of the sample using the technique.
INTRODUCTION
There are two general types of ion-exchange resins, the cation exchanger and the anion exchanger. Cation-
exchange resins contain acidic functional groups attached to the aromatic ring of the insoluble organic molecule.
Sulfonic acid groups, -SO3H, make up the active sites on the resin. Dowex 50 is a common and efficient cation-
exchange resin. After HCl treatment, the ionizable RSO3-H+ groups are formed on the resin. The hydrogen ions
exchange with other cations in the following manner:
Anion-exchange resins are composed of hydroxyl ions attached to basic group on the resin.
Going back to equation (12.1), since the exchange process is fast and complete, equivalents of displaced protons
can readily be determined by titration of the eluate (solution collected after separation) with standard NaOH.
Cation concentration is calculated from the amount of displaced hydrogen ions.
MATERIALS/APPARATUS
Absorbent cotton pH paper Iron stands
Burette clamps
CHEMICALS
NaOH HCl, concentrated Dowex 50 cation-exchange resin
Phenolphthalein
GLASSWARE
Burettes (50-mL) Volumetric pipettes (10-mL) Beakers (100-mL)
Erlenmeyer flasks (250-mL) Volumetric flasks (100-mL) Watch glass
PROCEDURE
NOTE: The cotton should not be very thick for easier removal of the resin and a more controllable flow
rate.
2. Fill up 1/4 of the ion-exchange column with the prepared resin. Gently pour the resin as it is suspended in
concentrated acid solution. Do not allow the level of liquid to fall below the resin level at any time. Wash
the sides of the column once some of the resin adheres to the sides. Open the column stopper to allow some
of the liquid to flow out of the column.
3. Wash out the excess acid in the column with distilled water until the pH of the eluate is equal to the pH of
the distilled water being used. Check pH using pH paper. Once pHeluate = pHdistilled water, the column is ready
for use.
NOTE: Test the pH of the latest eluate drop (by catching it in a watch glass) and not the pH of the eluate
in the receiving flask.
Storage of Resin
1. Pour 3.0 M HCl solution into the column until the blue color of the resin disappears and it reverts back to
its original yellowish-brown color.
2. To transfer the remaining resin to the original container, invert the burette and pour the resin with acid to
the original container. Repeat this step until all the resin has been transferred.
WASTE DISPOSAL
1. DO NOT dispose the excess resin. Return the resin to the instructor.
2. Drain all titrated solutions down the sink with copious amount of water.
3. Dispose of used cotton into the solid waste container.
Appendix 1
Laboratory Guidelines and Techniques in Analytical
Chemistry
LABORATORY GUIDELINES
1. All students are required to wear a lab gown and safety goggles during each experiment. This will be
strictly enforced to avoid accidents caused by chemical spills and other incidents.
2. Shorts, skirts, sandals, slippers are not allowed during experiments.
3. Avoid wearing contact lenses inside the laboratory. It sticks to the eyeball in the presence of organic
solvents.
4. Eating, drinking, smoking, and playing inside the laboratory are strictly prohibited.
5. All accidents, injuries, breakages and spills, no matter how minor, must be reported immediately to the
instructor.
6. Should a chemical get into your mouth, spit it out and rinse your mouth thoroughly with water. Similarly,
if any chemical comes into contact with any other parts of your body or clothes, wash thoroughly with
plenty of water.
7. Unauthorized experiments are strictly prohibited.
8. Unauthorized person(s) shall not be allowed in a laboratory.
9. The working area must be cleared of unnecessary materials. Put all bags and books in designated areas.
10. Do not bring reagent bottles to your working area.
11. Avoid wasteful use of reagents, water, and electrical power.
12. Solids, water, and other liquids spilled on your working area must be cleaned up as soon as possible.
13. Always pour waste reagents into their respective disposal jars (never in the sink, otherwise stated), as these
chemicals cause cumulative damage to our drainage system.
14. Deposit insoluble wastes such as paper, wood, glass, cork, etc. in the solid waste bin.
15. First aid kits and fire extinguishers are located in the respective preparation rooms.
16. Replace the top of every container immediately after removal of reagent.
17. Hold stoppers of reagent bottles between fingers. Never set a stopper on a desktop.
18. Never return any excess reagent to a bottle, unless specifically directed, to avoid contamination.
19. Avoid inserting spatulas into a bottle that contains a solid chemical. Instead, shake the capped bottle
vigorously to break up encrustation; then pour out desired quantity.
20. Procedures involving the liberation of volatile or toxic or flammable materials shall be performed in a fume
hood.
21. Before leaving see to it that:
a) your locker is locked,
b) your assigned working area is clean and dry, and
c) all floating equipment are returned to the instructor.
LABORATORY TECHNIQUES
Cleaning of glassware
1. Cleaning of glassware should be done prior to use.
2. Wash glassware with a detergent solution and then rinse initially with copious amounts of tap water and,
afterwards, with several portions of distilled water.
3. It is not necessary to dry the interior surface of glassware before use. It can cause contamination.
4. Never use a test tube brush when cleaning volumetric glassware.
5. Always rinse with distilled water after rinsing with copious amounts of tap water.
6. For pipettes:
a) Draw detergent solution to a level 2 to 3 cm above calibration mark using an aspirator.
b) Drain solution and then rinse pipet with several portions of tap water.
c) Fill pipet with distilled water, approximately 1/3 of its capacity and rotate so that entire interior is
wetted. Repeat this process twice.
7. For burettes:
a) Soak in a liquid detergent solution.
b) Drain solution and then rinse with several portions of tap water.
c) Fill burette with distilled water, approximately 1/3 of its capacity and rotate so that entire interior is
wetted. Repeat this process twice.
Aliquot Measurement
1. Rinse pipet with solution to be used before measuring out aliquot.
2. Forefinger must be faintly moist to facilitate easy control. Too much moisture makes control impossible.
3. Rinse pipet thoroughly after use.
4. Residual liquid is never blown out of a volumetric pipet or from some measuring pipets.
Dilution of Solutions
1. Solute should be dissolved completely before diluting to the mark.
2. Bring the liquid level almost to the mark and allow time for drainage.
3. Use a dropper/wash bottle to make final additions of solvent as are necessary.
4. Firmly stopper flask and invert it repeatedly to ensure thorough mixing.
5. Always add concentrated acid to water; never water to acid when diluting acid solutions.
Filling of Burette
1. Ensure that stopcock is closed.
2. Add 5 to10 mL of the titrant and rotate the burette to wet the interior completely. Allow the liquid to drain
through the tip. Repeat this at least two more times.
3. Fill the burette until above the zero mark.
4. Free the tip of air bubbles by rapidly rotating the stopcock and permitting small quantities of the titrant to
pass.
5. Lower the level of the liquid just to or somewhat below the zero mark.
6. Do not store base solutions in a burette for a long time, it can cause glass stopcocks to freeze upon long
contact.
Titration
1. Right-handed persons should use their right hand in swirling the flask and left hand in controlling the
stopcock.
2. Ensure that the tip of the burette is well within the titration flask.
3. Introduce the titrant in increments of about 1 mL and swirl constantly to ensure thorough mixing.
4. Decrease the size of the increments as the titration progresses. Addition should be dropwise when
endpoint is near.
5. Allow for drainage for at least thirty seconds before recording the final volume.
Reading of Volumes
1. Always read the lower meniscus which is the curved surface of a liquid at its interface with the atmosphere.
2. Read at eye level to avoid the apparent displacement of a liquid level as an observer changes position or
parallax error. Volume will appear smaller if read below eye level. Volume will appear larger if read above
eye level.
Weighing of Objects
1. Always allow an object that has been heated to return to room temperature before weighing since it causes
apparent weight of object to be low.
2. Use crucible tongs to prevent moisture uptake by dried objects during weighing.
3. Keep laboratory balance clean and neat. Clean up any spillages immediately.
4. Use an analytical balance for weighing solids to the nearest 0.1 mg or 0.0001 g, particularly in weighing
primary standards. The top loading balance can be used for weighing hygroscopic solids, such as sodium
hydroxide and potassium permanganate.
Filtration of Precipitate
Three steps are involved in filtering an analytical precipitate: decantation, washing, and transfer.
1. In decantation, as much supernatant liquid as possible is passed through the filter while the precipitated
solid is kept essentially undisturbed in the beaker where it was formed. This procedure speeds the overall
filtration rate by delaying the time at which the pores of the filtering medium become clogged with
precipitate.
2. Wash liquid is next added to the beaker and thoroughly mixed with the precipitate. The solid is allowed
to settle and the supernate is again decanted. Several washings may be required depending on the
precipitate. Most washing should be carried out before the solid is transferred.
3. In the transfer process, the bulk of the precipitate is moved from beaker to filter by suitably directed
streams of wash liquid.
Appendix 2
Instructions on Proper Use of Instruments
ANALYTICAL BALANCE
The analytical balances housed at the UP Diliman Institute of Chemistry (e.g. Mettler AE 200 Toledo) are high
precision weighing instruments. To maintain the quality performance of the balance, the user must know the
proper way of using the balance.
General Guidelines
1. Make sure that the area around the balance is clean and free of objects or materials not needed for the
weighing procedure.
2. The balance pan and the balance floor should be free of any dust or foreign matter. Protect the pan by
using a pan cover. Use camel’s hair brush for cleaning. Keep the doors of the balance chamber always
closed. It should only be opened when placing or removing the object to be weighed.
3. The balance should be levelled by means of a built-in spirit level before attempting to do any weighing
operation. The balance is equipped with a level indicator on the floor of the weighing chamber and two
adjustable levelling feet at the rear. Adjust the levelling feet until the bubble appears in the center circle of
the level indicator.
4. The following switches are often seen in analytical balances:
a) ON TARE – turns on the balance if it is off; zeros the balance
b) OFF – turns the balance off
c) MODE – selects weighing units, functions, or options
d) PRINT – sends weight data, statistical data, GLP data to computer or printer
For weighing purposes, one will be using only the ON TARE switch. Please do not press the other switches.
5. The balance should be allowed at least 20 minutes prior to use for warm up.
Operating Procedures
1. Press ON TARE to re-zero the display.
2. Open the door of the balance chamber and place the object or material to be weighed on to the pan. Use a
pan protector if available.
3. Wait for the stability indicator to appear before recording the weight of the object or material. The stability
indicator is the letter “S” that appears on the left side of the display window.
4. The full capacity of an analytical balance is 210 g, therefore do not weigh any object or material that is
greater than 210 g or you will destroy the balance. The bars on the upper right hand side of the display
window are the capacity bars. These bars indicate the percentage of the current weight to the balance
capacity.
Weighing by Addition
This is done by first determining the accurate mass of a dried container. Then, the desired quantity of the sample
is added to the vessel and increase in the mass is taken as the mass of the sample. For example,
1. Addition of the sample into the container is done by the use of a spatula or by tapping the sample vial such
that only small portions are added at one time until the desired mass is obtained. However, since the
analytical balance has a TARE button, when weighing objects or materials that must be held in a container,
pressing TARE may be done to store the weight of the container in the balance’s memory.
2. Press ON TARE to re-zero the display.
3. Open the door of the balance chamber and place the empty container to be weighed on to the pan. Its
weight is displayed and the stability indicator also appears.
4. Press ON TARE, the display blanks until stable weight readings are received, then indicates zero. The
container’s weight is stored-in memory.
5. Add material to the container. As the material is added, its net weight is displayed. When the stability
indicator appears, record the weight of the material.
6. Removing the container and material from the pan will cause the balance to display the container’s weight
as a negative number.
7. Press ON TARE to reset the balance to zero.
Weighing by Difference
This technique is especially useful when a series of samples of about similar size are to be weighed. It requires
only (n+1) weighings to obtain n samples, compared to 2n weighings for the addition method. In contrast to
weighing by addition, the receiving vessels need not be dried. This method is best suited when the sample to
be weighed should be protected from undue exposure to the atmosphere as in the case of a hygroscopic
material.
The weighing bottle containing the sample is weighed accurately and the balance reading zeroed. Then, an
approximate quantity of the sample is transferred into the receiving flask by tapping the weighing bottle slowly.
DO NOT use a spatula. The weighing bottle with sample is weighed again and the amount of the sample
transferred is checked as a negative value on the balance’s readout. If this is within the range for the sample
size, then the mass is recorded and the process is repeated for the next sample. If it is too little, the mass is
ignored and more of the sample is added into the receiving flask. Thus, the decrease in mass is taken as the
mass of the sample:
UV-VIS SPECTROPHOTOMETER
The UV-Vis spectrophotometer to be used for Chemistry 26.1 is Shimadzu UVmini-1240 single beam
spectrophotometer, housed at the Analytical Services Laboratory of UP Diliman Institute of Chemistry.
Operating Procedures
1. Turn on the spectrophotometer.
2. Determine analytical wavelength.
a) Fill up the cuvette with the standard solution with highest concentration (after rinsing the cuvette
properly with distilled water and solution to be measured).
b) Place the cuvette at the cuvette holder. Make sure the clear side faces the light source.
c) Go to Spectrum Mode.
d) Set the wavelength range.
e) Scan.
f) Obtain peak. This corresponds to the analytical wavelength.
3. Autozero.
a) Fill up the cuvette with blank or reference solution (after rinsing the cuvette properly with distilled
water and solution to be measured).
b) Place the cuvette at the cuvette holder. Make sure the clear side faces the light source.
c) Press Autozero for background correction.
4. Read absorbance measurements of the standard and sample solutions.
a) Press Go to WL and input the obtained analytical wavelength.
b) Starting from the standard solution with lowest concentration to highest concentration, fill up the
cuvette with the solution to be measured.
c) Place the cuvette at the cuvette holder. Make sure the clear side faces the light source.
d) Scan each solution.
e) Record the absorbance readings for each solution.
5. After reading all the standard and sample solutions, rinse all cuvettes with distilled water and allow it to
drain. Return the cells in the rack or in their storage container. Check with your instructor to see whether
the instrument is to be left on for other students; if not, turn it off.
Operating Procedures
1. Plug the pH meter into the power supply.
2. Select the pH mode on the function switch. Allow 15-20 minutes warm up.
3. Calibrate the pH meter before use:
a. Rinse the electrode with distilled water from a wash bottle. Gently blot off excess water with tissue
paper. This will minimize carryover and contamination.
b. Immerse the electrode in the pH 7 buffer solution. Allow sufficient time for the pH reading to
stabilize. Set the display to the correct value of the buffer using the BUFFER control.
c. Rinse the electrode with distilled water and blot dry with tissue paper.
d. Immerse the electrode in pH 4 or 10 buffer (the former if acidic substances will be measured, the
latter if basic solutions). Allow sufficient time for the pH reading to stabilize. Set the display to the
correct value of the buffer using the SLOPE control.
4. Measure the pH of the sample(s):
a. Rinse the electrode with distilled water and blot dry with tissue paper.
b. Immerse the electrode in the unknown solution. Allow the reading to stabilize. The display will
indicate the value of the solution directly in pH. Note the reading.
c. Rinse the electrode with distilled water and blot dry with clean tissue prior to immersing in the next
sample.
General Guidelines
1. Rinse the electrode thoroughly after use.
2. Handle the electrodes with care. Do not touch the sensitive glass pH membrane. Do not rub the electrode
as this may induce an electrostatic charge.
3. During use, ensure the electrode is rinsed between each measurement to eliminate contamination of
solutions.
4. Turn off pH meter after use.
Appendix 3
Preparation of Buffer Solutions
100 mL of 0.10 M HOAc-NaOAc buffer solution of pH 4.0 from 1.0 M HOAc and 1.0 M NaOAc
1. Buffer system:
2. Calculate the total moles of buffer components based of the given stock solutions to be used.
Total moles buffer = volume of buffer x conc. of buffer
= moles acid component + moles base component
= (0.100 L)(0.10 M) = 0.0100 mol
3. Using the Henderson-Hasselbalch equation, determine the number of moles of the acid and the base to
be used in preparing the buffer.
[𝑏𝑎𝑠𝑒]
𝑝𝐻 = 𝑝𝐾𝑎 + log
[𝑎𝑐𝑖𝑑]
[𝑏𝑎𝑠𝑒]
4.0 = 4.74 + log
[𝑎𝑐𝑖𝑑]
[𝑏𝑎𝑠𝑒]
log = −0.74
[𝑎𝑐𝑖𝑑]
[𝑏𝑎𝑠𝑒]
= 0.182
[𝑎𝑐𝑖𝑑]
(𝑀 𝑏𝑎𝑠𝑒)(𝐿 𝑏𝑎𝑠𝑒)
𝑡𝑜𝑡𝑎𝑙 𝑣𝑜𝑙 𝑏𝑢𝑓𝑓𝑒𝑟 𝑚𝑜𝑙𝑒𝑠 𝑏𝑎𝑠𝑒
= = 0.182
(𝑀 𝑎𝑐𝑖𝑑)(𝐿 𝑎𝑐𝑖𝑑) 𝑚𝑜𝑙𝑒𝑠 𝑎𝑐𝑖𝑑
𝑡𝑜𝑡𝑎𝑙 𝑣𝑜𝑙 𝑏𝑢𝑓𝑓𝑒𝑟
There are two working equations and two unknowns used to calculate for the moles of buffer
components.
𝑚𝑜𝑙𝑒𝑠 𝑏𝑎𝑠𝑒 = 0.182(𝑚𝑜𝑙𝑒𝑠 𝑎𝑐𝑖𝑑)
𝑚𝑜𝑙𝑒𝑠 𝑏𝑎𝑠𝑒 + 𝑚𝑜𝑙𝑒𝑠 𝑎𝑐𝑖𝑑 = 0.0100 𝑚𝑜𝑙
𝑚𝑜𝑙𝑒𝑠 𝑎𝑐𝑖𝑑 = 0.00846 𝑚𝑜𝑙
𝑚𝑜𝑙𝑒𝑠 𝑏𝑎𝑠𝑒 = 0.00154 𝑚𝑜𝑙
4. From the corresponding stock solutions, the volume of acid and base components can be calculated.
0.00846 𝑚𝑜𝑙
𝑣𝑜𝑙 1.0 𝑀 𝐻𝑂𝐴𝑐 = = 0.00846 𝐿 = 8.46 𝑚𝐿
1.0 𝑀 𝐻𝑂𝐴𝑐
0.00154 𝑚𝑜𝑙
𝑣𝑜𝑙 1.0 𝑀 𝑁𝑎𝑂𝐴𝑐 = = 0.00154 𝐿 = 1.54 𝑚𝐿
1.0 𝑀 𝑁𝑎𝑂𝐴𝑐
1. Buffer system:
2. Calculate the total moles of buffer components based of the given stock solutions to be used.
Total moles buffer = volume of buffer x conc. of buffer
= moles acid component + moles base component
= (0.100 L)(0.10 M) = 0.0100 mol
3. Using the Henderson-Hasselbalch equation, determine the number of moles of the acid and the base to
be used in preparing the buffer.
[𝑏𝑎𝑠𝑒]
𝑝𝐻 = 𝑝𝐾𝑎 + log
[𝑎𝑐𝑖𝑑]
[𝑏𝑎𝑠𝑒]
10.0 = 9.26 + log
[𝑎𝑐𝑖𝑑]
[𝑏𝑎𝑠𝑒]
log = 0.74
[𝑎𝑐𝑖𝑑]
[𝑏𝑎𝑠𝑒]
= 5.495
[𝑎𝑐𝑖𝑑]
(𝑀 𝑏𝑎𝑠𝑒)(𝐿 𝑏𝑎𝑠𝑒)
𝑡𝑜𝑡𝑎𝑙 𝑣𝑜𝑙 𝑏𝑢𝑓𝑓𝑒𝑟 𝑚𝑜𝑙𝑒𝑠 𝑏𝑎𝑠𝑒
= = 5.495
(𝑀 𝑎𝑐𝑖𝑑)(𝐿 𝑎𝑐𝑖𝑑) 𝑚𝑜𝑙𝑒𝑠 𝑎𝑐𝑖𝑑
𝑡𝑜𝑡𝑎𝑙 𝑣𝑜𝑙 𝑏𝑢𝑓𝑓𝑒𝑟
There are two working equations and two unknowns used to calculate for the moles of buffer
components.
𝑚𝑜𝑙𝑒𝑠 𝑏𝑎𝑠𝑒 = 5.495(𝑚𝑜𝑙𝑒𝑠 𝑎𝑐𝑖𝑑)
𝑚𝑜𝑙𝑒𝑠 𝑏𝑎𝑠𝑒 + 𝑚𝑜𝑙𝑒𝑠 𝑎𝑐𝑖𝑑 = 0.0100 𝑚𝑜𝑙
𝑚𝑜𝑙𝑒𝑠 𝑏𝑎𝑠𝑒 = 0.00846 𝑚𝑜𝑙
𝑚𝑜𝑙𝑒𝑠 𝑎𝑐𝑖𝑑 = 0.00154 𝑚𝑜𝑙
4. From the corresponding stock solutions, the volume of acid and base components can be calculated.
0.00846 𝑚𝑜𝑙
𝑣𝑜𝑙 𝑐𝑜𝑛𝑐 𝑁𝐻3 = = 0.000572 𝐿 = 0.572 𝑚𝐿
14.8 𝑀 𝑁𝐻3
53.49 𝑔
𝑚𝑎𝑠𝑠 𝑁𝐻4 𝐶𝑙 = 0.00154 𝑚𝑜𝑙 ( ) = 0.0824 𝑔
𝑚𝑜𝑙
5. Weigh out 0.0824 g NH4Cl crystals to the nearest 0.1 mg and add 0.572 mL of concentrated NH3.
6. Bulk the solution to the mark with distilled water and cover. Mix the solution thoroughly by repeated
shaking and inversion of the flask.
7. Determine the pH of the solution with a pH meter.
8. Adjust the pH if necessary by the addition of either an acid or base.
Appendix 4
Properties of Common Acids, Bases, and Primary
Standards
Table A4-1. Properties of common acid and base solutions.
Appendix 5
Tolerances of Common Laboratory Glassware and
Equipment
Table A5-1. Tolerances of common laboratory glassware and equipment.
25 mL 0.05 mL
Volumetric Pipettes
5 mL 0.01 mL Delivery of single and
10 mL 0.02 mL fixed volume between 0.5
25 mL 0.03 mL and 200 mL
50 mL 0.05 mL
Burettes
50 mL (Class A) 0.05 mL Titrations
50 mL (Class B) 0.10 mL
Volumetric Flasks If to contain If to deliver
50 mL 0.05 mL 0.10 mL
100 mL 0.08 mL 0.15 mL
Solution preparations
250 mL 0.12 mL 0.20 mL
500 mL 0.20 mL 0.30 mL
1000 mL 0.30 mL 0.50 mL
Analytical Balance 0.0002 g
Mass measurements
Top Loading Balance 0.01 g
Appendix 6
Significant Figures and Error Propagation
SIGNIFICANT FIGURES
The number of significant figures in a number consists of all of the certain digits and the first uncertain digit.
Rules
1. A zero may or may not be significant, depending upon its location in a number. A zero that is
surrounded by other digits is always significant because it is read directly and with certainty from a scale
or instrument readout.
Examples: 2.034 cm 4 significant figures
4.17 g 3 significant figures
4.0017 m 5 significant figures
2. Zeros that locate the decimal point for us are not significant.
Examples: 0.0034 g 2 significant figures
30.26 mL 4 significant figures
0.03026 L 4 significant figures
In an antilogarithm of a number, keep as many digits as there are digits to the right of the decimal point
in the original number.
Examples: antilog 12.5 = 3 x 1012
antilog -3.47 = 3.4 x 10-4
antilog 0.99 = 9.8
Rounding Numbers
1. In rounding a number ending in 5, always round so that the result ends with an even number. This
eliminates any tendency to round in a set direction.
Examples: 61.555 round to 4 significant figures 61.56
61.565 round to 4 significant figures 61.56
ERROR PROPAGATION
Addition
𝑅 ± 𝑟 = (𝐴 ± 𝑎) + (𝐵 ± 𝑏) + (𝐶 ± 𝑐)
𝑟 = √𝑎 2 + 𝑏 2 + 𝑐 2
Subtraction
𝑅 ± 𝑟 = (𝐴 ± 𝑎) − (𝐵 ± 𝑏) − (𝐶 ± 𝑐)
𝑟 = √𝑎 2 + 𝑏 2 + 𝑐 2
Multiplication
𝑅 ± 𝑟 = (𝐴 ± 𝑎) × (𝐵 ± 𝑏) × (𝐶 ± 𝑐)
𝑎 2 𝑏 2 𝑐 2
𝑟 = 𝑅 × √( ) + ( ) + ( )
𝐴 𝐵 𝐶
Division
𝑅 ± 𝑟 = (𝐴 ± 𝑎) ÷ (𝐵 ± 𝑏) ÷ (𝐶 ± 𝑐)
𝑎 2 𝑏 2 𝑐 2
𝑟 = 𝑅 × √( ) + ( ) + ( )
𝐴 𝐵 𝐶
Exponents
𝑅 ± 𝑟 = (𝐴 ± 𝑎) 𝑥
𝑎
𝑟 =𝑥×𝑅×
𝐴
Logarithm
𝑅 ± 𝑟 = log10 (𝐴 ± 𝑎)
𝑎
𝑟 = 0.434 ×
𝐴
Antilogarithm
𝑅 ± 𝑟 = antilog10 (𝐴 ± 𝑎)
𝑎
𝑟 = 2.303 × 𝑅 ×
𝐴
Appendix 7
Periodic Table of Elements