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.