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VPB 112 2009-10 1
Exercise 1
GENERAL INFORMATION ABOUT LABORATORY GLASSWARE
All laboratory work involves some form of measurement. Various types of volumetric
glassware are available measuring lab reagents. Before starting any biochemical experiment,
you must familiarize yourself with the types of volumetric glassware and should learn to use
the apparatus accurately to eliminate the errors resulting from careless handling of the
apparatus.
Volumetric glassware
A. Types of pipettes: The pipettes are designated class A or B according to their accuracy.
The class A pipettes are more accurate and conformity-certified. The class B pipettes are less
accurate but suitable for routine use.
1. Volumetric pipettes: These are one mark pipettes and are designated to deliver fixed
volumes. Before using the pipette it is rinsed with the solution to be pipetted, then filled to
just above the mark and after this the liquid is allowed to fall up to the mark. The tip of the
pipette is carefully wiped with tissue paper and the solution is allowed to drain into the
required container with tip of the pipette touching the wall of the container. After the solution
has drained, the tip of the pipette is touched with wall of the container for 15 seconds, then
removed. The solution present at the tip should not to be blown out.
2. Graduated pipettes: These pipettes are available in two types i.e. for total delivery or
blowout type, which are graduated to the tip. The interval between the calibration marks
depends upon the volume of the pipette. These pipettes are commonly used to measure out
odd quantities of liquid. It is better to use a 0.2 mL graduated pipette to deliver 0.1 mL rather
than a 5.0 mL pipette.
B. Burettes: These are graduated and used in volumetric titrations. Burettes deliver odd
quantities of liquid accurately. Burettes are available in 1, 2, 5, 10, 25 and 50 mL capacity.
The burettes designed to deliver up to 5 mL are known as micro burettes. These burettes have
narrow bore and liquid should be given time to drain to the required mark before taking the
reading. As with the pipettes, any liquid remaining at the tip should be removed by touching
the tip gently against the vessel. Burettes can be used for measuring poisonous or corrosive
liquids also.
C. Measuring cylinders: Measuring cylinder does not deliver the stated volume like a
pipette or a burette but only measures the liquid. Measuring cylinder can however, be used to
deliver relatively large volumes when exactness is not important.
D. Volumetric flasks: The volumetric flask has a narrow neck and is fitted with the cut glass
stopper. These flasks are calibrated to contain the volume specified at a fixed temperature.
VPB 112 2009-10 2
The volumetric flask is used for accurate measurement and adjustment of the volume of the
solution to be prepared. Before making up to the mark with solvent the solute is dissolved
first. If heating is required to dissolve the solute then suspension is transferred to the beaker
for heating and then cooled to room temperature and then transferred to the volumetric flask.
Never heat a volumetric flask, so it must not be dried in the oven because it will change its
volume.
Cleaning of glassware: The glassware used for laboratory work should be thoroughly
cleaned. Used glassware is washed in warm water containing detergent then rinsed several
times with tap water and then with glass distilled water. Do not use excess of detergent since
this may interfere with some of the experiments. Greasy glassware is cleaned by using
chloroform or toluene and then by soaking over night in chromic acid. Very dirty glassware
can be cleaned by soaking in a mixture of concentrated sulfuric acid and nitric acid if chromic
acid fails to clean the glassware. All traces of acid are removed by thoroughly rinsing in tap
water and then with glass distilled water. Normal glassware is dried in an oven, but
volumetric glassware is not heat dried but rinsed with little alcohol, then with ether and then
dried by blowing hot air.
Composition of chromic acid: Dissolve about 10 grams of sodium dichromate crystals in a
minimum quantity of water and add about 500 mL concentrated sulfuric acid.
VPB 112 2009-10 3
Exercise 2
PREPARATION OF AQUEOUS SOLUTIONS
A true solution is a homogenous mixture of one or more substances (solute) in a solvent
whose concentrations may be varied between certain definite limits. Most of the reactions
studied by biochemists occur in solutions having water as solvent and are termed as aqueous
solutions. Consequently, you should be familiar with various ways of expressing
concentrations of solutions and converting from one to the other.
A. Concentration based on volume: Concentration based on the amount of dissolved solute
per unit volume is the most widely used in various biochemical studies. The most common
conventions are as follows:
a) Molarity
b) Normality
c) Osmolarity
a) Molarity (M) = the number of moles of solute per liter of solution. Molar concentrations
are usually depicted in square brackets. e.g. [H+] = molarity of H ions.
Molarity( M) can be calculated from weight of dissolved solute and its MW. Wt g/MW =
Moles
To prepare a molar solution e.g. 100 mL of 0.1 M NaOH can be prepared by employing the
following calculations:
Liters x M = Number of moles of NaOH required or 0.1 x 0.1 = 0.01 moles of NaOH
required.
Now Number of moles = Wtg / MW,
therefore, Wtg = number of moles x MW
= 0.01 x 40 = 0.4g
Dilute solutions are expressed in terms of millimolar, micromolar, nanomolar or picomolar
etc.
Therefore, 1 mM = 10-3
M = 1 mmole/liter = 1 µmole/mL.
1 µM = 10-6
M = 1 µmole/liter = 1 nmole/mL.
1 nM = 10-9
M = 1 nmole/liter = 1 pmole/mL
A 1M solution contains one Avogadro number of molecules per liter and is frequently called
a mole regardless of nature of substance.
b) Normality (N) = the number of equivalents of solute per liter of solution.
N can be calculated from weight of the dissolved solute and its equivalent weight (EW)
i.e. Wtg / EW = Equivalents.
EW of an acid or base can be calculated from the number (n) of replaceable H+ or OH
- ions
per molecule.
EW= MW/n
VPB 112 2009-10 4
Another term milli-equivalent weight is often used and it is the weight in grams of acid or
base in 1 milliliter of 1 N solution. This term is employed to express the concentration of
electrolytes.
For conversion of mg % to milliequivalent (mEq) / liter, the following formula is used:
mEq / L= mg% x 10 / milliequivalent weight,
where mEq Wt. = millimolecular weight/valence
Normality and molarity are related by the expression, N = nM
e.g.0.IM solution of H2SO4 is 0.2 N
c) Osmolarity
Osmolarity of 1M solution of an ideal, non-dissociable solute is one.
Osmolarity = ni x molarity of particles in solution,
where ni = number of ions produced per molecule.
A 1M solution of a dissociable salt is ni x 1 = ni osmolar.
Similarly, a 0.2 M solution of NaCI is ni x 0.2 = 2 x 0.2 = 0.4 osmolar.
B. Concentrations based on weight
a) Molality
b) Mole fraction
a) Molality (m) = the number of the moles of solute per 1000 g of solvent. This mode of
measurement is used in the calculations of boiling point elevation and freezing point
depression. For dilute aqueous solutions, m and M are quite close.
b) Mole fraction- It is defined as the ratio of the number of moles of a compound present to
the total number of moles in the solution. For example, if a solution contains n1 moles of
compound 1, n2 moles of compound 2 and n3 moles of compound 3, the mole fraction of
compound 1, will be MF1 = nl/n1+n2 +n3.
MF x 100 = moIe percent.
C. Per cent solutions
a) Weight/weight per cent
b)Weight/volume per cent
c) Volume/volume per cent
a) Weight/weight per cent- It is expressed as weight in grams of a solute per 100 g of
solution. It is useful in the preparation of molar solution of commercial acids because their
concentrations are expressed in % w/w, if Specific Gravity of acid is known.
b) Weight/volume per cent- It is used in routine laboratory solutions where solute is a solid
while solvent is liquid and exact concentrations are not that important.
Weight/volume percent (% w/v) = mass of solute per 100 mL of solution.
In routine, the term gram per cent is used which refers to grams of solute per 100 mL (or dL-
deciliter) of the solution. Another term, milligram per cent is also used frequently in clinical
measurements: mg % = mg of solute per 100mL of solution.
c) Volume / volume per cent-It is used when the solute is also a liquid and a high degree of
accuracy is not required. For example, 70% (v/v) ethanol solution can be prepared by diluting
70 mL of absolute alcohol to 100 mL with water.
VPB 112 2009-10 5
Exercise 3
Preparation of a standard solution
Standard solutions
A standard solution is one whose exact strength or concentration is known and it is stable.
These are made either by weighing out a known mass of a primary standard substance,
dissolving it in the solvent, usually distilled water, and making up the solution to a specified
volume (e.g. sodium carbonate), or by standardization (after making up approximately) by
titration against a known standard (e.g. sulphuric acid against sodium carbonate). It is often
more convenient to prepare a solution a little more concentrated than required and then
diluting it with distilled water until the desired strength is achieved.
Primary standard substances:
A primary standard substance should satisfy the following requirements:
i) It must be easy to obtain, purify, dry (at 110-120°C) and preserve in pure state.
ii) It should not be hygroscopic.
iii) The substance should not have impurities exceeding 0.01-0.02 percent.
iv) It should have high equivalent weight so that the weighing errors may be negligible.
(For accuracy of 1 part in 1000, it is necessary to use samples weighing at least 0.2 g).
v) The substance should be readily soluble under the experimental conditions.
vi) The titration error should be negligible.
An alternative approach is to buy small quantities of standard solutions (ready-made or
concentrated) and use these to standardize one's own solutions. Once a solution has been
standardized so that its concentration is known precisely, it can be used as a 'secondary
standard' to standardize other solutions. So, a standard solution of sodium carbonate can be
used to standardize hydrochloric or sulphuric acids, which in turn can be used to standardize
sodium hydroxide solutions.
The substances commonly employed as primary standards are:
For acid base reactions: sodium carbonate (Na2CO3), borax (Na2B407), potassium hydrogen
pthalate KH(C8H404), potassium bi-iodate KH(IO3)2 , succinic acid 2(C4H404), benzoic acid
H(C7H5O2).
For precipitation reactions: - silver nitrate, sodium chloride, potassium chloride etc.
Secondary standard substances
Hydrated salts do not make good standards. However, some salts such as borax
(Na2B40710H20), oxalic acid (H2C2O4.2H2O) and copper sulfate (CuSO4.H2O) are found to be
satisfactory secondary standards.
VPB 112 2009-10 6
Preparation of 0.1N sodium carbonate solution (1L):
Materials:
1) Pure sodium carbonate - Analytical grade Na2CO3 is dehydrated by heating at 260-270°C
for 30 min and allowed to cool in a dessicator before use.
2) Distilled water
3) Balance
4) Volumetric flask 1 litre.
Procedure: Calculate the amount of sodium carbonate required to prepare 1 litre of 0.1 N
solution. Weigh out accurately the calculated amount of pure dry sodium carbonate in a clean
dry beaker, dissolve in about 200 mL of distilled water and transfer it to a 1 litre volumetric
flask. Give at least three washings to the beaker with DW and transfer these to the volumetric
flask. Make up the volume to the 1 litre mark with distilled water. Mix. This is 0.1 N solution
of sodium carbonate. Transfer to a reagent bottle and label it.
Preparation of 0.1N oxalic acid solution (1L):
Materials:
1) Pure, dry oxalic acid -Analytical grade
2) Distilled water
3) Balance
4) Volumetric flask 1 litre.
Procedure: Calculate the amount of oxalic acid required to prepare 1 litre of 0.1N solution.
Weigh out accurately pure oxalic acid in a clean dry beaker, dissolve in about 200 mL of
distilled water and transfer it to a 1 liter volumetric flask. Give at least three washings to the
beaker with DW and transfer these to the volumetric flask. Make up the volume to the 1 litre
mark with distilled water. Mix. This is 0.1 N solution of oxalic acid. Transfer to a reagent
bottle and label it.
VPB 112 2009-10 7
Exercise 4
PREPARATION AND STANDARDIZATION OF AQUEOUS SOLUTION OF ACIDS
Theory of neutralization indicators
The objective of titrating, say an alkaline solution with standard solution of an acid is the
determination of the amount of acid which is exactly equivalent chemically to the amount of
base present. The point at which it is reached is the equivalence point, stoichiometrical point,
or theoretical end point; an aqueous solution of the corresponding salt results. If both the acid
and base are strong electrolytes, the resultant solution will be neutral and have a pH of 7.
A large number of substances are available, called neutralization or acid-base indicators that
possess different colour according to the H+ concentration of the solution. The chief
characteristic of these indicators is that the change from predominantly acid colour to
predominantly alkaline colour is not sudden and abrupt, but takes place within a small
interval of pH (usually about two pH units). It is termed as the colour-change interval of the
indicator. The position of the colour change interval in the pH scale varies widely with
different indicators. For most acid-base titrations, that indicator can be selected which
exhibits distinct colour change at a pH close to that obtained at the equivalence point.
Preparation of a standard acid:
Two acids, namely, hydrochloric acid and sulfuric acid, are widely employed in the
preparation of standard solution of acids. Both of these are commercially available as
concentrated solutions; concentrated hydrochloric acid is about 10.5 -12N and concentrated
sulfuric acid is about 36 N. By suitable dilution, solutions of any desired approximate
strength can be prepared. Hydrochloric acid is generally preferred, since most chlorides are
soluble in water. Sulfuric acid forms insoluble salts with lime and baryta (BaOH2). For
titration of hot liquids or for determinations that require boiling for some time with excess of
acid, standard sulfuric acid is, however, prefered. Nitric acid is rarely employed, because it
almost invariably contains a little nitrous acid, which has a destructive action upon many
indicators.
Preparation of 0.l N Hydrochloric acid:
Materials:
1. Concentrated Hydrochloric acid
2. Distilled water
3. Graduated pipettes (5, 10 mL )
4. Volumetricflasks of required capacity
Procedure: Calculate the volume of conc. HCl required to prepare the required volume of
solution. Round off the calculated volume to next mL. Measure out this volume of
concentrated hydrochloric acid by means of a graduated pipette and pour the acid into a
volumetric flask half full of distilled water. (add acid to water not vice versa) Make up to the
mark with distilled water and thoroughly mix by shaking. This will give a solution
approximately 0.1N HCl.
VPB 112 2009-10 8
Standardization of the HCl solution with Sodium Carbonate
Materials:
1. Hydrochloric acid 0.1N (approx.)
2. Standard solution of Na2CO3 (0.1 N).
3. Methyl orange indicator.
4. Burette 25 mL.
5. Pipette 5 mL.
Procedure: Pipette 5 mL of standard solution of sodium carbonate (0.1 N) into a titration
flask. Add 1 drop of methyl orange indicator. Rinse a clean burette with solution of
hydrochloric acid already prepared. Fill the burette to a point 2-3 cm above the zero mark and
open the stopcock momentarily in order to fill the jet with liquid. Examine the jet to see that
no air bubbles are enclosed. If there are, more liquid must be run out until the jet is
completely filled. Note the initial reading of the burette. Place the conical flask containing
sodium carbonate solution upon a piece of unglazed white paper beneath the burette, and run
in the acid slowly from the burette. During the addition of the acid, the flask must be
constantly rotated with one hand whilst the other controls the stopcock. Continue to add the
acid solution drop wise until the contents in the flask becomes a faint red. This marks the end
point of the titration. Note the final reading of the burette accurately and determine the
volume used. Repeat the procedure until two concordant results are obtained.
Calculation of Normality:
If Vb is the volume in mL of the standard solution (Na2CO3) of normality Nb required to react
completely with Va mL of unknown solution (HCl) of normality Na
Then applying the normality equation,
Standard base Approx. Acid
Vb X Nb = Va X Na
or Na = Vb X Nb
Va
In this case Vb = 5 mL, Nb = 0.1, and suppose Va = x mL Then Na = 5 x 0.1/x = 0.5/x
From this solution of HCl having a initial normality Ni you want to prepare a standard HCl
solution of final normality Nf = 0.1 of known volume (Vf = say 100 mL). Applying the
normality equation again,
Initial acid Final acid
Vi X Ni = Vf X Nf
or Vi = Vf X Nf / Ni
Putting the values, Vi = 100 x 0.1 / 0.5/x = 20x mL
Measure out a volume equal to 20 x mL of the approx. 0.1 N HCl solution, dilute it to exactly
100 mL, mix and label it as 0.1 N HCl solution.
Verify your solution by titrating this standardized solution against the standard Na2CO3
solution. Note the volume of HCl solution used now. If equal volume of HCl solution is used
for titration of a known volume of Na2CO3, it confirms that normality of former is equal to
that of the latter.
VPB 112 2009-10 9
Exercise 5
Preparation and standardization of an Alkali Solution
The hydroxides of sodium, potassium and barium are generally employed for the preparation
of standard alkalis. They are strong bases and are readily soluble in water. Both potassium
and sodium hydroxides are extremely hygroscopic; a certain amount of alkali carbonate and
water are always present. Exact results cannot be obtained in the presence of carbonate with
some indicators; therefore, carbonate free alkali solutions are prepared.
Preparation of approximately 0.1N sodium hydroxide
Materials:
1. Sodium hydroxide -A.R. grade (carbonate free).
2. Distilled water
3. Balance
4. Volumetric flask of required capacity
Procedure: Calculate the amount of NaOH required and round it off to next gram. Rapidly
weigh it on a watch glass or in a small beaker, dissolve in a smaller volume of DW, make up
to the required volume with boiled out distilled water, mix thoroughly by shaking and pour
the resultant solution into a reagent bottle, which should be closed by a rubber stopper. This
solution is labeled as approximately 0.1 N NaOH solution.
Standardisation with standard oxalic acid
Materials:
1 Approximately 0.1N NaOH
2. Standard solution of oxalic acid 0.1N
3. Phenolphthalein indicator
4. Burette 25 mL
5. Pipette 5 mL
Procedure: Place the approx. 0.1N NaOH solution in the burette. Transfer 5 mL of the 0.1 N
oxalic acid solution into a 100 mL conical flask with the help of a pipette and add 1-2 drops
of phenolphthalein indicator and titrate with sodium hydroxide solution from the burette.
Appearance of a faint pink color that stays for at least 30 seconds marks the end point. Repeat
the procedure until two concordant results are obtained.
General Calculations:
If VA is the volume in mL of the standard solution (Oxalic acid) of normality NA required to
react completely with VB mL of unknown solution (NaOH) of normality NB
Then applying the normality equation,
Standard Acid Approx. Base
VA X NA = VB X NB
or NB = VA X NA
VB
In this case VA = 5 mL, NA = 0.1, and suppose VB = x mL
So NB = 5 x 0.1/x = 0.5/x
VPB 112 2009-10 10
From this solution of NaOH having a initial normality Ni you want to prepare a standard
NaOH solution of final normality Nf = 0.1 of known volume (Vf = say 50 mL).
Applying the normality equation again,
Initial base Final base
Vi X Ni = Vf X Nf
or Vi = Vf X Nf / Ni
Putting the values,
Vi = 50 x 0.1 / 0.5/x
= 10x mL
Measure out a volume equal to 10 x mL of the approx. 0.1 N NaOH solution, dilute it to
exactly 50 mL, mix and label it as 0.1 N NaOH solution.
Verify your solution by titrating this standardized solution against the standard oxalic acid
solution. Note the volume of NaOH solution used now.
Questions :
1. What is molarity?
2. What is the difference between a molar and a molal solution?
3. What is normality?
4. How can a 0.9 % solution of NaCl be prepared?
5. What is osmolarity?
6. What is a primary standard?
7. How strength of a solution can be calculated?
8. How mg/dl can be converted to mmol/L?
9. Why phenolphthalein is not used as an indicator while standardizing HCI with Na2CO3?
VPB 112 2009-10 11
Exercise 6
DETERMINATION OF pH
The hydrogen ion concentration of most solutions and body fluids is extremely low. In 1909,
Sorenson introduced the term pH as a convenient way of expressing hydrogen ion
concentration, which avoids the use of cumbersome numbers. pH is defined as the negative
logarithm of hydrogen ion concentration (To be exact it should be hydrogen ion activity, but
in dilute solutions the hydrogen ion concentration is virtually the same as the activity).
pH = - log [H+]
e.g. plasma [H+] = 0.00000025 M = 2.5 x 10
-7mol/L
or plasma pH = - log (2.5 x 10-7
) = 7.4
pH of the solution can be determined by colorimetric and electrometric methods.
A. Determination of pH by colorimetric methods:
An approximate determination of the pH of a liquid may be made by the use of indicator test
papers or by the systematic use of number of indicators or by the use of multiple range
indicator solution (universal indicator).
i) Litmus paper
The use of the litmus paper can simply tell if the solution being tested is acidic or basic. The
blue litmus changes to red in acidic solution while the red litmus changes to blue if it is basic.
ii) Indicator Test Paper Method
Material:
1. Wide range test paper (pH paper) covering the pH ranges 1-14, 2-10 etc.
2. Narrow range test paper (pH paper) covering pH ranges 1-5, 5-9 etc.
3. Colour matching charts.
4. Samples to be tested.
Procedure: Pour the sample into a small beaker and dip the test paper into the sample.
Compare the colour of the test paper with that of the colour-matching chart that shows the
change in colour at regular pH intervals.
Limitations: 1.The test papers tend to deteriorate upon storage.
2. For average observer the test papers do not permit the determination of
pH closer than 0.5-1 pH units.
VPB 112 2009-10 12
iii) Buffer Solution Method:
pH indicators are organic compounds of natural or synthetic origin whose colour is dependent
upon the pH of the solution. Indicators are usually weak acids that dissociate in solution.
Indicator = Indicator - + H
+
Applying the Henderson-Hasselbalch equation,
pH = pKIn + log10( Indicator -) / (Indicator)
The two forms of the indicator have different colors and, as can be seen from above equation,
the actual colour of the solution will depend upon the pKIn and the pH. The greatest colour
change occurs around the pKIn and this is where the indicator is most useful.
A series of appropriate buffer solutions is selected, differing successively in pH by about 0.2,
covering the pH range of the solutions under investigation. The range of the buffer solutions
required is indicated by the preliminary pH determination.
Materials:
1. Standard buffer solutions differing successively in pH by 0.2 units (e.g. Citrate buffer, pH
3.0 -6.2, phosphate buffer, pH 5.8 -8.0, barbital buffer, pH 6.8 -9.2 and Tris buffer, pH 7.2 -
9.0).
Phosphate buffer (0.1M, pH 5.8 -8.0). It can be prepared as follows.
i) 0.2 M Na2HPO4 - dissolve 28.4 g of Na2HPO4 in distilled water and make the volume up
to one litre.
ii) 0.2 M NaH2PO4 -dissolve 31.2 g of NaH2PO4 .2H2O in distilled water and make the
volume up to one litre.
Prepare the following buffer solutions by mixing the above solutions.
Na2HPO4 (0.2M),
mL
NaH2PO4 (0.2M),
mL
Final Volume
mL
pH
8.0 92.0 200 5.8
12.3 87.7 200 6.0
18.5 81.5 200 6.2
26.5 73.5 200 6.4
37.5 62.5 200 6.6
49.0 51.0 200 6.8
61.0 39.0 200 7.0
72.0 28.0 200 7.2
81.0 19.0 200 7.4
87.0 13.0 200 7.6
91.5 8.5 200 7.8
94.7 5.3 200 8.0
2. Samples to be tested (Blood, cerebrospinal fluid, saliva, milk, rumen liquor, urine etc.)
diluted 1 in 10.
VPB 112 2009-10 13
3. Indicators (List of indicators given in Table given below).
Table: Colour change and useful pH range of some common indicators.
Indicator/colour Acid Base pKa Useful pH
Cresol red Red Yellow - 0.2 – 1.8
Thymol blue Red Yellow 1.7 1.2 – 2.8
Metacresol
purple
Red Yellow 1.5 1.2 – 2.8
Bromophenol
blue
Yellow Blue 4.0 3.0 – 5.0
Methyl orange Red Orange 3.7 3.1 – 4.4
Bromocresol
green
Yellow Blue 4.7 3.8 – 5.4
Methyl red Red Yellow 5.0 4.3 – 6.0
Bromocresol
purple
Yellow Purple 6.3 5.5 – 7.0
Bromothymol
blue
Yellow Blue 7.1 6.0 – 7.6
Phenol red Yellow Red 7.9 6.8 – 8.2
Phenolphthlein Colourless Pink 9.7 8.3 – 10.0
Procedure: Place equal volumes (say 10 mL) of the buffer solution differing successively in
pH by about 0.2, in test tubes of colourless glass and having approximately the same
dimensions. Add a small amount of a suitable indicator for the particular pH range to each
tube. A series of different colours corresponding to different pH values is obtained.
Treat an equal volume (say 10 mL) of the test solution with equal volume of indicator
to that used for the buffer solutions, and compare the resulting colour with that of the
coloured standard buffer solutions. When the complete match is found, the test solution and
corresponding buffer solution have the same pH. Sometimes a complete match is not
obtained, but the colour of the test solution falls between those of two successive standards.
Further buffer solution may then be prepared differing by 0.1 pH, if desired, and then pH
value re-determined. As a general rule colorimetric methods cannot be relied upon to give
values of pH more accurate than to within 0.2 pH unit.
For determination of the pH of a coloured solution, e.g. urine, the sample is treated
with indicator as previously outlined. The problem due to the yellow colour of the urine can
be solved by placing the tube containing only urine behind the buffer standard and the tube
containing only water behind the tube of urine plus indicator and observing the colour of light
passing through the pairs of tubes. In this way, the colour of the urine is made constant for
both the standard and unknown sample, and correct pH values may be obtained. Moreover,
pH indicators are affected by oxidizing agents, reducing agents, salt concentration and
protein, so these facts must be borne in mind when using them.
Precaution: Add only a small quantity (1-2 drops) of indicator to the solution under
examination.
VPB 112 2009-10 14
B. Determination of pH by electrometric method
The most convenient and reliable method for measuring pH is by the use of a pH meter which
measures the e.m.f. of a concentration cell formed from the reference electrode, the test
solution and a glass electrode sensitive to hydrogen ions.
Glass electrode
A glass electrode consists of a very thin bulb about 0.1 mm thick, made up of soft glass
which is pH-responsive e.g. corning 015 glass, containing 72% SiO2, 22% Na2O2 and 6%
CaO. This pH responsive glass bulb is sealed to a stem of harder high-resistance glass. Inside
the bulb is a solution of HCI (0.1 mol/L) connected to a platinum wire via a silver-silver
chloride electrode. A potential develops across the thin glass of the bulb which depends on
the pH of the solution in which it is immersed. This potential is not readily affected by salts,
proteins, or oxidizing and reducing agents. So the electrode can be used in a wide variety of
media. The glass electrode in the test solution constitutes a half cell and the measuring circuit
is completed by a reference electrode which is not sensitive to hydrogen ions.
Calomel electrode
The reference electrode commonly used is the calomel electrode similar to that illustrated. It
consists of a solution of KCI in contact with a solution of Hg2CI2 (calomel) and mercury. The
calomel electrode is stable, easily prepared, and the potential with respect to the standard
hydrogen electrode is accurately known. The reference electrode must be in contact with the
test solution via a liquid junction which generally is KCI that diffuses slowly out of electrode
to give electrical continuity.
Combination electrode
The reference electrode is built around the glass electrode in a combination electrode. It
consists of an Ag / AgCI internal electrode and external reference electrode.
pH Meter:
The e.m.f. of the complete cell (E) formed by linking of these two electrodes is
E = Eref - Eglass
Where, Eref is the potential of calomel reference electrode (which at normal room temperature
is + 0.250V) and Eg1ass is the potential of the glass electrode which depends on the pH of the
solution under test.
Materials:
1. pH meter.
2. Standard buffers of 4.0, 7.0 and 9.2 pH for the calibration of pH meter.
3. Solutions to be tested.
VPB 112 2009-10 15
Procedure: For the measurement of pH with the help of a pH meter, the stepwise procedure
is as follows:
1. Keep the selector in standby position.
2. Switch on the pH meter and wait for 15 minutes. The digital display/needle should indicate
0.00.
3. Wash the electrode with distilled water, wipe it clean and dip it in the standard buffer
solution whose pH is close to that under test. Set the temperature of the pH meter at the
temperature of the standard buffer solution.
4. Bring the selector to pH mode and calibrate the instrument so that display should indicate
the pH of the standard buffer. Then bring the selector back to standby position.
5. Remove the electrode from the standard buffer solution, wash with distilled water and store
it in buffer solution (pH 4.0) until further use.
6. To measure the pH of the test solution/sample, adjust the temperature of the pH meter to
the sample temperature. Then dip the electrode in the test solution, bring the selector to pH
mode and read the pH from the digital display.
7. After measurement of pH, bring the selector back to standby position. Wash the electrode
thoroughly with distilled water particularly after measuring the pH of a solution with high
concentration of biological macromolecules, as these may adhere to the glass and distort
subsequent pH measurements unless removed. After use, the electrode is stored in buffer
solution (pH 4.0) and must never be allowed to dry out.
Normal pH values of various body fluids.
Body fluid pH Body fluid pH
Aqueous humor of eye 7.4 Liver (intracellular) 6.4 – 7.4
Blood serum/plasma 7.35 – 7.45 Milk 6.6 – 6.9
Cerebrospinal fluid 7.35 – 7.45 Pancreatic juice 7.0 – 8.0
Feces 7.0 – 7.5 Rumen liquor 5.0 – 7.0
Bladder bile 6.0 – 7.7 Saliva
(monogastric)
6.35 – 8.35
Gastric juice 1.0 – 3.0 Saliva (ruminants) 8.0
Hepatic bile 7.1 – 7.3 Tears 7.0
Intestinal juice 7.0 – 8.0 Urine (omnivorous) 4.8 – 7.5
Intestinal fluid 7.4 Urine (carnivorous) 4.0 – 5.0
Urine (herbivorous) 7.0 – 8.0
VPB 112 2009-10 16
Exercise 7
PREPARATION OF A BUFFER OF KNOWN pH
A pH buffer is a substance or mixture of substances, which permits solutions to resist large
changes in pH upon addition of small amounts of H+ or OH
- ions. Common buffer mixtures
contain two substances, a conjugate acid and a conjugate base. An acidic buffer contains a
weak acid and a salt of the weak acid (conjugate base). A basic buffer contains a weak base
and a salt of the weak base (conjugate acid). Together the two species (conjugate acid plus
conjugate base) resist large changes in pH by partially absorbing additions of H+ or OH
- ions
to the solution.
To prepare a buffer solution of known pH e.g. 7.0, a weak acid whose pKa is 7.0 or close to
7.0 is used. Henderson-Hasselbalch equation is used to calculate the molar ratio of the
conjugate base and conjugate acid at the given pH and pKa of the weak acid.
Example:
For the preparation of 2 liters of a 1M buffer, pH = 7.0 and pKa = 7.0, proceed as follows:
pH = pK + log[CB] /[ CA] CB = Conjugate Base, CA = Conjugate Acid
7.0 = 7.0 + log [CB] /[ CA]
0.0 = log [CB] /[ CA]
or [CB] /[ CA] = Antilog 0.00 = 1 = 1/1
1/2 of total = CB and 1/2 of total = CA
1/2 x 1M = 0.5 Therefore, CB = 0.5M
1/2 x 1M = 0.5 Therefore, CA = 0.5 M
For 2 liters of 1M buffer; 2 liters x 1M = moles of total (CB & CA)
The total of 2.0 moles is obtained from two sources;
2 liters x 0.5M = 1 mole CB and 2 liters x 0.5M = 1 mole CA
Now number of moles of CB = Wtg CB/ MWCB
or WtgCB = 1 x MW CB
, say X g
Similarly WtgCA = 1 x MW CA
, say Y g
To prepare the buffer solution, dissolve X g of conjugate base and Y g of conjugate acid in
DW, make up the volume to near the final volume, check pH and then make up the required
volume with DW.
VPB 112 2009-10 17
Prepare 1 liter of 0.2 M acetate buffer having a pH 5.0.
Materials:
1. pH meter
2. Standard buffer solutions for calibration
3. CH3COONa
4. Acetic acid glacial.
Procedure: Calculate the weight of conjugate base and conjugate acid using Henderson-
Hasselbalch equation. Dissolve the required amounts of conjugate base and conjugate acid in
distilled water. If the total volume required is 1 liter, make the volume up to approximately
900 mL, check the pH with the help of pH meter and then make the volume up to 1 liter with
distilled water.
pK values of some Acids and Bases useful in preparing Buffers
Free acid or base pK at 250C
Citric acid 3.06 (pK1)
Acetic acid 4.76
Carbonic acid 6.1 (pK1)
Phosphoric acid 7.21 (pK2)
Tris 8.1
Carbonic acid 10.25 (pK2)
Experiment 1: Prepare 1 L of 0.2 M acetate buffer, pH 4.46 from solid NaCOOCH3.3H2O
and 1M solution of acetic acid.
Experiment 2: Prepare 1 L of 0.2 M phosphate buffer, pH 7.0 from Na2HPO4 and
NaH2PO4.2H2O.
Experiment 3: Prepare 1 L of 0.2M carbonate buffer, pH 10.7 from Na2CO3 and NaHCO3.
Questions:
1. What is a buffer?
2. What is buffering capacity and what are the factors that affect it?
3. What is pK value?
4. What are physiological buffers and what is their importance?
VPB 112 2009-10 18
Exercise 8
Qualitative Analysis of Carbohydrates
1. Solubility: The monosaccharides and oligosaccharides are readily soluble in water due to
polar hydroxyl groups, which forms H-bonds with water. The polysaccharides owing to their
large molecular weight, however, make translucent colloidal solutions.
2. Qualitative tests for Carbohydrates: While analyzing a sample containing a mixture of
carbohydrates, particularly the sugars, several difficulties are encountered in their qualitative
as well as quantitative analysis. These difficulties are attributed to their structural and
chemical similarity and also with respect to their stereoisomerism. Therefore, during
biochemical investigation it becomes necessary to establish whether a given sample contains
carbohydrates or not. Several rapid tests are available to establish the presence or absence of
a sugar or a carbohydrate in a sample. These tests are based on specific colour reactions
typical for their group. In the laboratory, it is advisable to perform these tests with the
individual rather than mixture of sugars. The sensitivity of these tests can be confirmed by
using sugar solutions of different concentrations (0.1- 1%).
A. General tests for carbohydrates: The most commonly used tests to detect the presence
of carbohydrates in a solution are:
a) Molisch’s test: It is a group test for all carbohydrates, whether free or in combined form.
Despite its limitations, it is routinely used to detect the presence of carbohydrates.
Principle: The reaction is based on the fact that concentrated H2SO4 catalyses the dehydration
of sugars to form furfural (from pentoses) or hydroxymethyl furfural (from hexoses).These
furfurals then condense with sulfonated alpha-naphthol to give a purple or violet coloured
product. Polysaccharides and glycoproteins also give a +ve reaction. In the event of the
carbohydrate being a poly- or disaccharide, the acid first hydrolyses it into component
monosaccharides, which then get dehydrated to form furfural or its derivatives.
Reagents: i) Conc.H2SO4
ii) Molisch’s reagent: Alpha-naphthol 5%(w/v) in 95% ethanol.
Procedure: Take 1-2 mL of unknown solution and add 2-3 drops of Molisch’s reagent and
mix the contents. Incline the tube and carefully pour 1-2 mL of conc.H2SO4 down the side of
VPB 112 2009-10 19
tube so that the acid forms a layer beneath the aqueous solution. The formation of a purple or
violet ring or zone at the junction of two layers indicates the presence of carbohydrates.
Precautions: i) Alpha-naphthol solution is unstable and should be prepared fresh.
ii) The conc.H2SO4 should be added carefully along the sides of the test
tube causing minimal disturbance to the contents of the tube.
Limitations: - In addition to carbohydrates, furfurals as such, some organic acids, aldehydes
and ketones also give this test. Secondly, a concentrated sugar solution may give a red colour
instead of purple owing to charring action of acid.
b) Anthrone test:
Principle;- Anthrone reaction is another general test for carbohydrates. Its principle is same as
that for Molisch’s test except that the furfurals and hydroxy-methyl furfurals give
condensation products with anthrone that are bluish green in colour.
Reagents: i) Anthrone reagent: 0.2%(w/v) solution in conc.H2SO4.
Procedure: Add about 2 mL of Anthrone reagent to about 0.5-1mL of the test solution in a
test tube and mix thoroughly. Observe whether the colour changes to bluish green. If not,
then examine the tubes again after keeping them in boiling water bath for ten minutes. A blue
green colour indicates positive test.
B. Specific tests for carbohydrates:
a) Iodine test for polysaccharides: This test is performed to distinguish polysaccharides
from mono- and disaccharides.
Principle: Iodine forms coloured adsorption complexes with different polysaccharides.
These complexes are formed due to the adsorption of iodine on the polysaccharide chains.
The intensity of the colour depends on the length of the unbranched or linear chain available
for the complex formation. Thus, amylose, the unbranched helical component of starch gives
a deep blue colour and amylopectin, the branched component gives red colour because the
chains do not coil effectively. Glycogen, which is also highly branched, gives red colour with
iodine. This test is conducted in acidic or neutral solutions.
VPB 112 2009-10 20
Reagents: i) Iodine solution: Prepare 2%(w/v) solution of KI in water to which add a few
crystals of iodine until the solution assume a deep yellow colour.
ii) Starch solution: Dissolve 1g starch in about 10-20mL boiling water and
make the volume to 100mL with saturated sodium chloride solution.
Procedure: Take 2-3 mL of the test solution in a test tube and add 1-2 drops of dil.HCl.
Mix and then add 1-2 drops of iodine solution. Mix and observe the colour change. Heat the
tube and observe the colour again. Blue colour disappears on heating and reappears on
cooling.
b) Tests based on reducing property of carbohydrates: Sugars possessing a free, or
potentially free, aldehyde or ketone group act as reducing agents and this fact becomes the
basis of the tests performed for distinguishing them from the non reducing sugars. Such
sugars have the property of readily reducing alkaline solutions of the metals like copper,
bismuth, mercury, iron and silver. The aldo sugars are oxidized to the corresponding aldonic
acids whereas the keto sugars give rise to shorter chain acids. If the alkaline copper solution
is heated in the absence of reducing sugar, it forms black precipitate of cupric oxide:
Heat Cu (OH)2 ------------→ CuO + H2O In the presence of a reducing sugar, however, the alkaline solution of copper is reduced to
insoluble yellow or red cuprous oxide:
Heat
Sugar + 2 Cu(OH)2 ------------→ Aldonic acid + Cu2O + 2 H2O
i) Fehling’s test: Rochelle salt acts as chelating agent in this reaction:
CuSO4 + 2KOH ------------→ Cu(OH)2 + K2SO4
2Cu(OH)2 + Reducing Sugar ------------→ 2Cu2O + Aldonic acid
Reagents: i) Fehling’s solution A: Dissolve 69.38 g of Copper Sulfate in DW and make
the volume to 1 L.
ii) Fehling’s solution B: Dissolve 250 g NaOH in DW, add 346 g of Sodium
Potassium Tartrate and make the volume up to 1 L.
Mix equal volumes of A & B solutions just before use because mixing
causes deterioration with time.
Procedure: Add 1mL of Fehling’s reagent to 1mL of the test solution. Mix thoroughly and
Place the test tubes in boiling water bath. Formation of yellow or red precipitates of Cuprous
Oxide indicates the presence of reducing sugar.
Note: i) In case of mild reduction, leave the solution to stand for 10-15 minutes, then decant
the supernatant. A small amount of red or yellow precipitates may then be seen adhering to
the inner side of the tube.
ii) Fehling’s test is performed only alkaline solution.
iii) Cuprous Oxide is dissolved by ammonia. Hence it is not possible to detect small
quantities of reducing sugars in fluids saturated with ammonium salts e.g. urine.
ii) Benedict’s Test: Benedict modified the Fehling’s solution to produce an improved single
reagent which quite stable. Sodium Citrate functions as a chelating agent. It is very sensitive
and even small quantities of reducing sugars(0.1%) yield enough precipitates.
VPB 112 2009-10 21
Reaction:
Reagents: i) Benedict’s qualitative reagent: Dissolve 173g Sod. Citrate and 100g anhydrous
Sod. Carbonate in about 800mL water by gently heating the contents. Then in a separate
beaker dissolve 17.3g Copper Sulfate in about 100mL DW. Pour this solution slowly, with
constant stirring into the Carbonate-Citrate mixture and make upto 1 L with DW.
Procedure: Add 0.5-1mL of the test solution to about 2mL of Benedict’s reagent. Keep the
test tubes in boiling water bath. Observe the formation of green, orange, yellow or red
precipitates which indicates the presence of reducing sugar in the given solution.
Note: i) This test is especially suitable for the detection of reducing sugar in urine because it
is more specific than Fehling’s test which is also positive for non-reducing substances
such as urates present in urine.
ii) This is a semi-quantitative test.
iii) Barfoed’s Test: This test is performed to distinguish between a reducing mono- and
disaccharide. Monosaccharides are more reactive reducing agents than disaccharides and thus
react in about 1-2 min while the reducing disaccharides take 7-12 min to get hydrolysed in
the acidic solution and then react. Thus, the difference in reducing property can be detected.
Reaction:
Reagents: i) Barfoed’s reagent: Dissolve 66.5 g of Cupric Acetate in about 900 mL DW. Boil
and add 9 mL of Glacial Acetic Acid. Cool and make the volume to 1 L with DW and filter if
necessary.
Procedure: Take 2-3 mL of Barfoed’s reagent in a test tube and add 1 mL of the given test
solution. Keep the test tubes in boiling water bath for 1-2 min only. Then allow the tubes to
cool down for a while. Thin red precipitates, at the bottom or sides of the tube indicates the
presence of a reducing monosaccharide.
VPB 112 2009-10 22
Note: i) The boiling should not be prolonged beyond 1-2min, otherwise reducing
disaccharides also respond to this test.
ii) This test does not work in the detection of reducing sugar in urine owing to the
presence of chloride ions.
iv) Picric Acid Test: It is another test for the detection of reducing sugars. The reducing
sugars react with Picric Acid to form a red coloured Picramic Acid.
Reagents: i) Saturated picric acid: Dissolve 13 g Picric Acid in 100mL DW, boil and cool.
ii) Sodium Carbonate (10% solution).
Procedure: Add 1mL of the above reagent to 1mL of the test solution followed by 0.5 mL of
10% Sod. Carbonate solution. Heat the test tube in a boiling water bath. Appearance of red
colour indicates the presence of reducing sugar in the solution.
c) Seliwanoff’s test for keto sugars
Principle: This test is a timed colour reaction specific for keto hexoses. Thus it is used to
distinguish aldoses from ketoses. In the presence of HCl ketohexoses undergo dehydration to
yield 4-hydroxy methyl furfural more rapidly than aldohexoses. Further these furfural
derivatives condense with resorcinol to form a red coloured complex.
Reagents: i) Seliwanoff’s reagent: Disolve 50mg resorcinol in 100 mL dilute HCl (1:2).
Procedure: To about 2 mL of Seliwanoff’s reagent add 1 mL of the test solution and warm in
a boiling water bath for 1 min. Appearance of a red colour indicates the presence of
ketohexose (fructose).
Note: i) Aldohexoses e.g. glucose also react if boiling is prolonged because it is transformed
into fructose by the catalytic action of acid.
ii) Sucrose and inulin also give this test because these are hydrolysed by acid to
give fructose.
d) Bial’s test for pentoses
Principle: This test is specific for pentoses and the compounds containing pentoses and thus
useful for the determination of pentose sugars. Reaction is due to the formation of furfural in
the acid medium which condenses with orcinol in the presence of ferric ions to give a blue
green coloured complex.
VPB 112 2009-10 23
Reagents: i) Bial’s reagent: Dissolve 1.5 g of orcinol in 100 mL of conc. HCl and add 20-30
drops of 10% Ferric Chloride solution to it. Prepare fresh.
Procedure: To about 2 mL 0f Bial’s reagent add 4-5 drops of test solution. Heat in a boiling
water bath until bubbles of gas rise to the surface. Formation of green solution and precipitate
indicates the presence of a pentose sugar.
e) Test for sucrose: This test is performed only when there is no precipitation in Barfoed’s
test.
Principle: Sucrose present in the unknown solution is hydrolysed by acid to glucose and
fructose. The resulting fructose formed in the solution is then tested by Seliwanoff’s reagent.
Reagents: i) Conc. HCl
ii) Seliwanoff’s reagent
iii) Sodium Carbonate
Procedure: To about 2-3mL of the test solution add1-2 drops of conc. HCl and boil in a water
bath for about 8-10 minutes. Then add about 5mL of Seliwanoff’s reagent and again keep it
in the water bath for 1minute. Appearance of red colour indicates the presence of fructose
which is the hydrolytic product of sucrose.
Note: Acid hydrolyzed sample after cooling and then neutralizing with Sodium Carbonate
can be tested by Benedict’s reagent for reducing sugars.
f) Mucic acid test for galactose
Principle: This test is highly specific for galactose which is either independently present in
solutions or obtained by the hydrolysis of lactose. Galactose is converted to Saccharic
acid on heating with HNO3(a strong oxidizing agent). Mucic acid (galactaric acid) which is
formed from galactose due to the oxidation of both aldehyde & primary alcoholic group at
C1&C6. It is the only Saccharic acid which is insoluble in cold water and thus helps in the
identification of galactose.
Reagents: i) Conc. HNO3
Procedure: Take about 50mg galactose and 50mg glucose separately in test tubes. Add 1mL
DW and 1mL conc. HNO3 to each tube. Heat the tubes in a boiling water bath for about 1hr.
Add 5mL DW and let the tubes to stand and cool slowly. Colourless needle like crystals will
indicate the presence of galactose.
Note: Lactose will also give this test.
g) Phenylhydrazine test / Osazone Test
This test is used to differentiate the maltose and lactose
Principle:
An organic compound phenylhydrazine reacts with carbonyl carbon of sugar to form the
osazones. These osazone crystals have yellow colour characteristics shapes and melting point,
time of formation and solubility. The characterstics features of osazone are given in the
following table:-
VPB 112 2009-10 24
Carbohydrate
(Osazone)
Time of formation
(Minutes)
Solubility in boiling
water
Crystalline structure
Fructosazone 2 Insoluble Needle shape
Glucosazone 5 Insoluble Needle shape
Galactosazone 20 Insoluble Thorny ball shape
Maltosazone 30-45 soluble Sunflower/Star shape
Lactosazone 30-45 soluble Cotton ball/Powder puff shape
Procedure:
Take 7-8 ml of carbohydrate solution in a test tube and to this add a pinch of phenylhydrazine
and double the quantity of sodium acetate and 10 drops of acetic acid. Dissolve by shaking and
allow cooling slowly. Observe the shape of crystal under low power of microscope (10x).
Observations and inference:
The lactose forms powder puff shape crystals, maltose forms sunflower shaped or star shaped
crystals, while the glucose and fructose form identical needle shaped crystals.
VPB 112 2009-10 25
Exercise 9
QUALITATIVE TESTS FOR PROTEINS AND AMINO ACIDS
Proteins are high molecular weight, nitrogen containing organic compounds, composed of
amino acids linked through peptide bonds. They are found in every part of every cell, as they
are fundamental in all aspects of cell structure and function. The quantity and type of proteins
show a wide variation not only among different organisms but also in different organs or
parts of the same organism. A number of tests are performed in order to detect the presence
of proteins, identify the constituent amino acids, and then identify the proteins themselves,
on the basis of amino acid composition and certain other properties specific of a particular
protein.
Experiments on proteins and amino acids
A. Hydrolysis of proteins:
i) Hydrolysis with alkali:
Principle: Alkaline hydrolysis causes the racemization of naturally occurring levo form of
alpha amino acids in proteins, producing a racemic mixture of D & L forms. This hydrolysis
also causes partial or complete decomposition of cystine, arginine and lysine. The amide-
groups of glutamine and asparagines release ammonia on alkaline hydrolysis. Sulfur
containing amino acid, cystine is partially destroyed to release unoxidised sulfur in the
hydrolysate.
Reagents: 10N NaOH: Dissolve 40g of sodium hydroxide in distilled water and make the
volume to 100mL.
Procedure: Take about 2-3 mL of test solution in a test tube and add equal volume of 10N
NaOH solution. Observe the precipitation. Place the tubes in a boiling water bath and stir
occasionally. Observe the colour and odour of the hydrolysate. Then take 1 mL of the
hydrolysate and add 10 drops of lead acetate solution. A black precipitate is formed which
indicates the presence of unoxidised sulfur.
ii) Hydrolysis with Acid:
Principle: Acid hydrolysis does not affect the optical activity of the amino acids to any
appreciable extent, but there is complete destruction of tryptophan and results in the
formation of humin, a black pigment. Certain other amino acids e.g. methionine and tyrosine
are partially lost if metals are present. Glutamine and asparagines are deamidated.
Reagents: 6N H2SO4: Take 16.6mL of conc.H2SO4 and make the volume to 100mL with
distilled water.
Procedure: Take about 5 mL of egg albumin solution in a test tube and add 5mL of 6N
H2SO4. Observe the precipitates formed. Heat the tube in a boiling water bath, stir
occasionally and observe the change in colour and odour as the hydrolysis proceeds. Take 1
mL of the hydrolysate and add 10 drops of lead acetate solution. Compare the amounts of
black precipitates formed by lead acetate in two types of hydrolysis.
VPB 112 2009-10 26
B. Precipitation of protein with strong acids ( Heller’s Test )
Principle: Strong mineral acids like HNO3, HCl and H2SO4 cause denaturation of proteins
which results in their precipitation.
Reagent: Concentrated HNO3
Procedure: Place 1-2 mL of conc. Nitric acid in a test tube. Incline the tube and add an equal
volume of test solution. A white ring formed at the junction of two liquids indicates the
presence of protein.
Note: This test is used to detect albumin in urine. Other strong acids also give this test but
less readily than HNO3.
C. Tests for peptide linkages (two or more):
i) Biuret test:
Principle: When a solution of biuret reacts with sodium hydroxide and dilute cupric sulfate, a
purple violet colored co-ordination complex is formed. A co-ordination complex of same
chemical nature and color is also formed when proteins or peptides(except for a dipeptide)
react with alkaline cupric sulfate. Evidently, −CO−NH− groups joined either by a nitrogen (in
biuret) or a carbon (in protein and peptide) are responsible for this test and hence the name
biuret test, the reaction being positive for protein and peptides bearing at least two peptide
bonds. The co-ordination complex is formed between cupric ions and unshared electron pairs
of peptide nitrogen as shown below.
Reagents:
i) Biuret Reagent: Add, with stirring, 300 mL of 10% (w/v) NaOH to 500 mL of a solution
containing 0.3% copper sulfate pentahydrate and 1.2% sodium potassium tartrate, then dilute
to one liter. The reagent is stable for a few months but not a year. Adding one gram of
potassium iodide per liter and storing in the dark makes it stable indefinitely. The reagent can
be used for either qualitative or quantitative estimations.
ii) Test solution: 0.5% solution of a protein like casein or BSA in NaOH or egg albumin
solution.
VPB 112 2009-10 27
Procedure: Take about 1 mL of the test solution in a test tube and add 2-3 mL of biuret
reagent, mix thoroughly, and let stand for 10 minutes at room temperature. Appearance of a
purplish-violet or pinkish-violet color indicates the presence of protein. Repeat the test
substituting DW for test solution and compare the color.
Note: i. It is important to note that excess addition of cupric sulphate is not desirable as
it causes precipitation of copper hydroxide posing a difficulty in making out the
purple-violet color change.
ii. The shade of the colour produced depends upon the nature of the test protein,
being pink for peptones, blue for gelatin and purple-violet to pink-violet for
other proteins.
iii. Salts like ammonium sulfate and magnesium sulfate interfere with the test.
To avoid it the test solution should be diluted with two volumes of 40% NaOH.
D. Tests to study general properties of amino acids
i) Ninhydrin test:
Principle: When amino acid or proteins are heated with ninhydrin (triketo hydrindene
hydrate), they are oxidatively deaminated to give a reduced form of ninhydrin (hydrindantin),
CO2,NH3 and an aldehyde. The final purple coloured product, called Ruhman’s Purple
(diketo hydrindamine), is formed as a consequence of reaction between ninhydrin , ammonia
and hydrindantin. The reaction takes place under acidic conditions. Therefore, in case of
alkaline test solutions (eg. Casein, being soluble in alkali) a few drops of 1% acetic acid
should be added to ensure neutrality, before performing this test.
The imino acids proline & hydroxy proline, react with ninhydrin giving yellow coloured
product. This test is also given by certain amines and acid amides. The reaction is very
sensitive and suitable for the detection of minute quantities of amino acids.
Reagents: i) 0.2% solution of ninhydrin in water.
ii) 1% solution of amino acids/ proteins.
Procedure: Take 2-3 mL of a test solution in a test tube and add 5-6 drops of ninhydrin
solution. Heat in a boiling water bath for two minutes. A purple or blue colour, appearing on
cooling the contents, indicates the presence of a protein or amino acid. This test is positive for
all the members of protein family, namely proteins, peptones, peptides and amino acids.
VPB 112 2009-10 28
ii) Xanthoproteic test
Principle: This test is based on the ability of aromatic amino acids containing substituted
phenyl group, to react with conc. HNO3 to give dinitro derivatives of yellow colour. The salts
of these derivatives are, however, orange in colour. Tyrosine and tryptophan containing
proteins give a positive test, but phenylalanine does not respond to this test under ordinary
conditions, possibly due to the presence of unsubstituted phenyl groups in its structure.
Reagents: i) Conc. HNO3
ii) Sodium hydroxide solution, 40%(w/v).
Procedure: Take 2-3 mL of egg albumin in a test tube and add 1 mL of conc. HNO3. A white
precipitate is formed which on slight heating turns yellow and finally dissolve to impart
yellow colour to the solution. Cool and carefully add 2-3 mL of 40% NaOH solution and mix
with caution. The colour of the solution turns to deep yellow or orange, indicating the
presence of tyrosine and tryptophan amino acids in the protein.
iii) Millon- Nasse’s test:
Principle: This test is given by tyrosine, both free and constitutive of proteins.
Hydroxyphenyl group (C6H5OH) of tyrosine yields a purple red nitro-hydroxy-phenyl
mercurial, on heating with acid mercuric sulfate and sodium nitrite. Other compounds like
phenol, thymol, salicylic acid etc. also give this reaction.
Reagents: i) Acid mercuric sulfate - A 10% (w/v) solution in H2SO4.
ii) Sodium nitrite - A freshly prepared solution (1%w/v) in water.
Procedure: Add 1mL of acid mercuric sulfate solution to 2mL of protein solution taken in a
test tube. Mix and boil cautiously for 2 minutes. Cool under tap and add 1 mL of sodium
nitrite solution and warm gently. A deep red colour indicates the presence of tyrosine.
iv) Aldehyde test:
Principle: This test also specifically confirms the presence of tryptophan. Sulfuric acid in the
presence of mercuric sulfate, acting as an oxidizing agent, oxidizes indole ring of trytophane.
VPB 112 2009-10 29
This oxidized product of the indole nucleus reacts with formaldehyde (formaline) to give a
reddish violet complex.
Reagents:
i) Mercuric sulfate solution. It is the same reagent as used in Millon-Nasse’s
reaction.
ii) Dilute formaline. Take 2 mL of 40% formaline and dilute it to one litre with
water.
iii) Conc. Sulfuric acid.
Procedure: Add 1-2 drops of dilute formaline to 2-3 mL protein solution. Mix and add one
drop of mercuric sulfate solution. Mix again and incline the tube, and carefully add 2mL of
conc. Sulfuric acid in such a way that it runs down gently along the side of the tube to make
an acid layer beneath the protein mixture. Revert the test tube back to upright position and
gently rotate between the palms of your hands. A reddish violet colour appears at the zone of
contact of two liquids.
Note: Gelatin, a protein totally deficient in tryptophan, does not answer the above test. The
fact is taken as a confirmation of the presence of gelatin as no other general proteins are
deficient in tryptophan.
v) Tests for sulfur containing amino acids
a) Fohl’s reaction:
Principle: This test is performed to detect the presence of cysteine and cystine amino acids
containing weakly bound sulfur. It is based on the property of proteins, containing Cysteine
or cystine among its constituent amino acids, to yield sodium sulfide on heating in an alkaline
medium (provided by NaOH). The sodium sulfide thus formed reacts with lead acetate to
give lead sulfide. However, sulfur in methionine is too strongly bound to be released by
strong alkali and thus it does not respond to this test. Proteins like casein and gelatin, which
contain sufficient amounts of methionine but negligibly small amounts of either cysteine or
cystine, do not answer this test.
Reagents: i) Lead acetate- 2% solution(w/v) in water.
ii) Strong sodium hydroxide- 40%(w/v) in water.
Procedure: To 3 mL of protein solution, add an equal volume of strong sodium hydroxide.
Boil the contents for 1-2 minutes and then add 5-6 drops of lead acetate solution. The
solution darkens due to the formation of black precipitates of lead sulfide.
b) Nitroprusside test:
Principle: This test is also positive for cysteine and cystine amino acids. It is based on the fact
that sodium sulfide, produced by the alkali treatment of the test protein, forms a violet red
complex on reaction with sodium nitroprusside. Proteins casein and gelatin do not give this
test due to the negligibly small amounts of cysteine and cystine.
Reagents: i) Strong sodium hydroxide solution (40%w/v in water).
ii) Sodium nitroprusside solution (5%w/v in water).
VPB 112 2009-10 30
Procedure: Take 2-3 mL of protein solution and add an equal volume of strong sodium
hydroxide solution. Boil the contents for 3 minutes. Cool and add 3-4 drops of sodium
nitroprusside solution. A violet red coloration results, indicating the presence of cysteine and
cystine amino acids.
vi) Test for Arginine - Sakaguchi test:
Principle: This test detects the presence of arginine, an amino acid containing guanidino
group in its structure. An additive complex is formed between alpha-naphthol and guanidine-
group of arginine under alkaline conditions. This complex is oxidized by sodium
hypobromite (or sodium hypochlorite) to produce a bright red colour.
Reagents: i) Alpha-naphthol solution (1%w/v in 95 percent alcohol).
ii) Sodium hydroxide solution (10%w/v in water).
iii) Sodium hypobromite- In a 100 mL measuring cylinder, dissolve 40g of
sodium hydroxide in water and make a total volume of 75-80 mL. Add 10
mL of bromine to this solution cautiously with constant shaking and cooling.
Make a final volume of 100 mL by adding water.
Procedure: Take about 3 mL of the test solution and add 1 mL of 5% NaOH solution, 2 drops
of 1% alpha-naphthol solution followed by the addition of 10 drops of sodium hypobromite
solution. Mix well. If arginine is present, a bright red colour develops on standing.
Note: i) The reagents always give some colour. Therefore, a control should also be run.
ii) This test is very sensitive and can be used as a general test for all the proteins.
VPB 112 2009-10 31
Exercise 10
ESTIMATION OF AMINO ACID BY SORENSON’S TITRATION METHOD
Principle: Amino acids react with formalin to form methelyne amino acids,
H2NCH2CO2H + HCHO CH2=NCH2CO2H + H2O
But it was found that the reaction is more complex and the main product is monomethylol
and dimethyllol amino acids.
--NH2 + HCHO --NH(CH2OH)
-NH(CH2OH) + HCHO --N (CH2OH)2
Amino acids exist in zwitter ionic forms and cannot be titrated directly with alkali.
Thus, amino groups of amino acids are blocked by reaction with formaldehyde (or it reduces
the basic character of amino group). But formaldehyde does not react with the charged amino
groups (-NH+
3), thus, first the amino acid reacts with sodium hydroxide solution to give
NH2CHRCOO-, which condenses with formaldehyde to give a stable anion.
-OH + H3NCHRCOO
- → H2NCHRCOO
-+H2O
CH2O + H2NCHRCOO- → CH2 = NCHRCOO- + H2O
+ CH2OHNHCHRCOO-
+ (CH2OH)2NCH2COO-
Further, formaldehyde solutions contain formic acid and also amino acids are not
exactly neutral, thus it is necessary that both the formaldehyde and amino acid solutions
should have the same pH before mixing and for this purpose each solution is first made just
alkaline to phenolphthalein by means of dilute sodium hydroxide solution.
Reagents:
i) Neutral formalin solution: It is prepared by neutralization of formalin solution
with alkali to phenolphthalein end point.
ii) Standard glycine solution: Accurately weigh 0.8g glycine. Dissolve in distilled
water in a volumetric flask to100ml.
iii) Unknown glycine solution
iv) Sodium hydroxide solution (0.1 N)
v) Phenolphthalein indicator
Procedure: i) Transfer 20 mL of standard glycine solution in a conical flask and add 1-2 drops
of phenolphthalein.
VPB 112 2009-10 32
ii) Titrate with NaOH (0.1 N) to a faint pink color.
iii) Now add 10 mL (excess) of neutralized formaldehyde solution. The pink color of
the solution in the flask disappears.
iv) Titrate again with NaOH solution till the faint pink color is obtained.
v) Repeat the titration to get three concordant readings.
vi) Repeat the steps i) – v) for the unknown glycine solution.
Observations:
Weight of glycine dissolved in 100 mL of standard glycine solution = w g
Volume of NaOH solution used with 20 mL of standard glycine solution = v1 mL
Volume of NaOH solution used with 20 mL of unknown glycine solution = v2 mL
Calculations:
v1 mL of NaOH = 20 mL of standard glycine soln = w/5 g of glycine
v2 mL of NaOH = 20 mL of unknown glycine solution = w/5 . v2/v1 g of glycine
Thus strength of glycine in the unknown solution = w/5 . v2/v1 . 1000/20 g/L
= w.v2.10/v1 g/L
VPB 112 2009-10 33
Exercise No. 11
CHEMICAL CONSTANTS OF FATS
Some measurements, when undertaken on oils and fats convey valuable information
concerning their composition, purity, age and adequacy of storage conditions. These are
preferred over a complete chemical analysis of fats because such analysis involves long, time
consuming and cumbersome procedures.
a) Saponification value:
The triacylglycerols may be readily hydrolyzed to glycerol and salts of the constituent fatty
acids (soaps) by boiling with strong bases such as NaOH or KOH. The saponification value is
defined as “the number of milligrams of KOH required to saponify one gram of fat.” Since
fats are mixtures of glycerides containing different carbon chain length fatty acids, the
saponification number is an index of the molecular size of the fatty acid present in a fat.
b) Acid value:
Fat and oils contain some fatty acids in addition to the fatty acids, which are esterified to
glycerol. These free fatty acids can be estimated by titrating against an alkali. Thus, acid no.
can be defined as “the number of milligrams of KOH required to neutralize the free fatty acid
present in 1g of fat.” It is of value in determining rancidity due to free fatty acids. During
storage, fat may become rancid as a result of peroxide formation at the double bonds by
atmospheric oxygen and hydrolysis by micro-organisms with the liberation of free acids. The
amount of free fatty acids present gives an indication of the age and quality of fat.
c) Iodine value:
Iodine number may be defined as the number of grams of iodine absorbed by 100 g of oil or fat.
It indicates the degree of unsaturation of a fat or oil. Halogens add across the double bonds of
unsaturated fatty acids present in fat or oil to form additional compounds. Oils like soybean, corn
and cottonseed have high Iodine numbers because of the presence of more unsaturated fatty acids
in the fat molecules. Coconut oil, on the other hand, has a low Iodine value.
d) Reichert- Meissel Number : Fats and oils contain fatty acids of varying lengths and
R.M.number is a measure of the short chain fatty acids (volatile) present in fat. VFAs can be
obtained by saponification of fat, acidification to liberate the free fatty acids and then steam
distillation. These VFAs are then neutralized with standard alkali. Thus, R.M. number may be
VPB 112 2009-10 34
defined as the milliliters of 0.1N alkali required to neutralize the VFAs obtained from 5 g of
fat.
e) Acetyl Number: The acetyl number is defined as the number of milligrams of KOH
required to neutralize the acetic acid obtained by saponification of 1 g of fat after it has been
acetylated. (Acetylation refers to the treatment of fat or fatty acid mixture with acetic
anhydride which results in acetylation of all alcoholic OH groups). The acetyl number is,
thus, a measure of OH groups in a given sample of fat or oil. For example the castor oil has a
high acetyl number (146) because of high content of a hydroxy acid, ricnoleic acid, in it.
Butter, on the other hand, has a low acetyl value of 1.9 to 8.6, indicating the presence of a
small amount of hydroxylated acids.
Analytical constants of some common fats
Lipid Saponification
Number
Acid
Number
Iodine
Number
R.M.
number
Acetyl
Number
Lard 195 – 203 0.5 – 0.8 47 – 66 0.5 – 0.8 2.6
Human Fat 194 – 198 - 65 – 69 0.25 – 0.55 -
Tallow Beef 196 – 200 0.25 35 – 42 - 2.7 – 8.6
Butter Fat 210 – 230 0.4 26 – 28 17 – 35 1.9 – 8.6
Olive Oil 185 – 196 0.3 – 1.0 79 – 88 0.6 – 1.5 10 – 11
Cotton Seed
Oil
194 – 196 0.6 – 0.9 103 – 111 0.95 21 – 25
Linseed Oil 188 – 195 1.0 – 3.5 175 – 202 0.95 4.0
Castor Oil 175 – 183 0.12 – 0.8 84 1.4 146 – 150
Coconut Oil 253 – 262 1.1 – 1.9 6 – 10 6.6 – 7.5 2.0
VPB 112 2009-10 35
Exercise 12
DETERMINATION OF SAPONIFICATION VALUE OF FAT
Principle: The triacylglycerols can be hydrolyzed to glycerol and salts of the constituent
fatty acids (soaps) by boiling with strong bases such as NaOH or KOH. The saponification
value is defined as “the number of milligrams of KOH required to saponify one gram of fat.”
Since fats are mixtures of glycerides containing different carbon chain length fatty acids, the
saponification number is an index of the molecular size of the fatty acid present in a fat.
H2C-O-CO-C17H35
|
HC-O-C17H35
|
H2C-O-CO-C17H35
STEARIN
+ KOH →
H2COH
|
H2COH
|
H2COH
Glycerol
+ 3 C17H35COOK
Potassium Stearate (Soap)
From the above reaction, it may be noted that three molecules of KOH are consumed for
saponification of each molecule of triacylglycerol irrespective of chain length of fatty acid.
Evidently, each gram of a triacylglycerol with shorter chain fatty acids will contain larger
number of molecules of the triacylglycerol and thus requires much more KOH. The
saponification value is thus an indication of average molecular weight of the fatty acids in an
acylglycerol.
Reagents:
i) Standard HCL (0.5N).
ii) 0.5N alcoholic KOH - dissolve 28 g of KOH pellets in little water and make
the volume to 1 L with 95% ethyl alcohol.
iii) Phenolphthalein indicator
Procedure: Weigh accurately about 2 g of an oil or fat sample in a 250 mL conical flask.
Add 25 mL of 0.5 N alcoholic KOH to the flask. Similarly take 25 mL of 0.5 N alcoholic
KOH in another flask but without oil (blank). Fit each flask with reflux air condenser and
heat on a boiling water bath for 30 minutes. Cool the flask, add 1 drop of phenolphthalein
VPB 112 2009-10 36
indicator to each flask and titrate with 0.5 N HCL. Record the volume of HCL used for each
titration.
Observations & Calculations:
Weight of oil = A g
0.5N HCL used for sample titration = x mL
“ “ “ “ blank “ = y mL
volume of 0.5 N KOH used for saponification = (y-x) mL
Now 1 mL of 0.5 N HCL ≡ 1 mL of 0.5 N KOH ≡ 0.028 g of KOH
Saponification No. = (y-x) X 0.028 X 1000
A
VPB 112 2009-10 37
Exercise 13
DETERMINATION OF ACID VALUE OF FAT
Principle: Fat and oils contain some fatty acids in addition to the fatty acids that are
esterified to glycerol. These free fatty acids can be estimated by titrating against an alkali.
Thus, acid no. is defined as “the number of milligrams of KOH required to neutralize the free
fatty acid present in 1 g of fat”. The acid value is useful in determining rancidity due to free
fatty acids. During storage, fat may become rancid because of peroxide formation at the
double bonds by atmospheric oxygen and hydrolysis by microorganisms with the liberation
of free acids. Thus, the amount of free fatty acids present in a fat sample gives an indication
of the age and quality of fat.
Reagents: i) 95% ethyl alcohol.
ii) Phenolphthalein indicator
iii) 0.1N KOH solution-5.6g KOH/L
Procedure: Take about 5 mL of oil in a conical flask and weigh it accurately. Transfer 25
mL of 95% ethyl alcohol to another flask and add a drop of phenolphthalein indicator. Heat
the contents of flask to boiling and add 0.1 N KOH solution until a faint pink color remains.
Add this neutral alcohol sample to the flask containing oil sample, heat to boiling and titrate
to faint pink color with 0.1 N KOH.
Observations & Calculations:
Weight of oil = A g
Volume of 0.1 N KOH used for blank = x mL
“ “ “ “ “ “ sample = y mL
Actual volume of 0.1 N KOH used = y-x mL
Now 1 mL of 0.1N KOH = 5.6 mg of KOH
Acid no. = (y-x) X 5.6
A
VPB 112 2009-10 38
Exercise 14
Determination of Lactose in a sample of milk.
When an aqueous solution of picric acid containing glucose is heated with an alkali, there
results a Burgundy red color the intensity of which is dependent upon the quantity of sugar in
solution. On this reaction is based the Lewis-Benedict1 technique for the determination of
blood sugar. Lactose likewise produces a color with picric acid and alkali, but also that the
color is proportional to the amount of lactose present.
Determination of Lactose in Milk
Technique.
Standard.-Dissolve in a saturated aqueous solution of picric acid 0.1 per cent of. pure
recrystallized lactose. Each cc. of standard, therefore, is equivalent to 1 mg. of lactose.
Preparation of Milk: Dilute the milk to be tested 1 to 100. To about 10 cc. of the diluted milk
placed in a centrifuge tube, add about 0.5 gm. of solid picric acid, and dissolve by stirring and
shaking. Permit the solution to rest for about 5 minutes, then centrifuge at the usual rate
(1,500 revolutions) for 10 minutes. The result is a clear yellow liquid between a sediment of
precipitated proteins and a thin layer of fat. Filter through a small filter paper to remove the
fat. Into a long test-tube pipette 1 cc. of the milk solution. Into a separate tube introduce 1 cc.
of the standard. To each tube, add 1 mL of a saturated sodium carbonate solution. Mix and
immerse the tubes in boiling water for 20 minutes. Dilute, when cooled, both the standard and
the test solution to 10 cc. Read absorbance at nm
Calculation of Results.-To find the amount of lactose in the milk the following formula may
be used:
where x is the percentage of lactose in the milk, S the reading of the standard, R the reading
of the unknown, m the mg. of lactose to which the standard corresponds, and W is the
number of cc. of milk used. As a rule, S is set at 10, m’is 1, and W is 0.01.
A picramic acid standard may also be used.
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