A laboratory Text book of Biochemistry, Molecular Biology and Microbiology

CONTENTS

1. WATER, pH AND BUFFERS
1.1 PREPARATION OF ACETATE BUFFER
1.2 TESTING OF BUFFER CAPACITY

2. DILUTIONS, UNITS FOR MEASURING CONCENTRATION OF SOLUTIONS AND SPECTROPHOTOMETRIC CALCULATIONS

3. LAB SAFETY RULES

4. CARBOHYDRATES
4.1 QUALITATIVE TESTS FOR CAROHYDRATES
4.2 DETERMINATION OF REDUCING SUGARS CONTENT (NELSON-SOMOGYI METHOD)
4.3 DETERMINATION OF SUGARS CONTENT (DNSA METHOD)
4.4 DETERMINATION OF TOTAL CARBOHYDRATES (ANTHRONE METHOD)
4.5 DETERMINATION OF TOTAL CARBOHYDRATES (PHENOL SULFURIC ACID METHOD)
4.6 DETERMINATION OF GLUCOSE CONTENT (GLUCOSE OXIDASE METHOD)
4.7 DETERMINATION OF STARCH CONTENT
4.8 DETERMINATION OF CELLULOSE CONTENT
4.9 DETERMINATION OF FRUCTOSE CONTENT
4.10 DETERMINATION OF INULIN CONTENT

5. LIPIDS
5.1 QUALITATIVE TESTS FOR LIPIDS
5.2 QUALITATIVE TESTS FOR CHOLESTEROL
5.3DETERMINATION OF CHOLESTEROL CONTENT (ZAK'S METHOD)
5.4 DETERMINATION OF OIL/LIPID CONTENT IN OILSEEDS
5.5 DETERMINATION OF VOLATILE ACIDS CONTENT
5.6 DETERMINATION OF ACID VALUE AND % FREE FATTY ACIDS
5.7 DETERMINATION OF SAPONIFICATION NUMBER OF OIL/FAT
5.8 DETERMINATION OF IODINE NUMBEROF OIL/FAT
5.9 DETERMINATION OF PEROXIDE VALUE
5.10 SEPARATION OF PLANT PIGMENTS USING THIN LAYER CHROMATOGRAPHY

6. PROTEINS
6.1 QUALITATIVE TEST FOR AMINO ACIDS AND PROTEINS
6.2 PRECIPITATION OF PROTEINS (A GENERAL APPROACH)
6.3 DETERMINATION OF lMAX FOR A GIVEN PROTEIN SAMPLE
6.4 DETERMINATION OF PROTEIN CONTENT (LOWRY’S METHOD)
6.5 PRECIPITATION AND DETERMINATION OF PROTEIN CONTENT (TCA METHOD)
6.6 DETERMINATION OF PROTEIN CONTENT (BIURET METHOD)
6.7 DETERMINATION OF PROTEIN CONTENT (BRADFORD METHOD)
6.8 DETERMINATION OF PROLINE CONTENT IN PLANTS
6.9 ISOLATION AND PURIFICATION OF CASEIN FROM MILK
6.10 DETERMINATION OF TOTAL FREE AMINO ACIDS
6.11 DETERMINATION OF TRYPTOPHAN CONTENT
6.12 DETERMINATION OF METHIONINE CONTENT
6.13 DETERMINATION OF LYSINE CONTENT
6.14 SEPARATION OF AMINO ACIDS BY CIRCULAR PAPER CHROMATOGRAPHY
6.15 SEPARATION OF AMINO ACIDS USING PAPER CHROMATOGRAPHY
6.16 SDS-PAGE
6.17 PREPARATION OF STANDARD CURVE OF PARA NITRO PHENOL (PNP)
6.18 ENZYMATIC ASSAY OF ACID PHOSPHATASE
6.19 DETERMINATION OF THE TIME LINEARITY OF AN ENZYMATIC REACTION
6.20 DETERMINATION OF THE PROTEIN LINEARITY OF AN ENZYMATIC REACTION
6.21 TO DETERMINE THE PROPERTIES OF (KM AND VMAX VALUES) OF AN ENZYME
6.22 NITRATE REDUCTASE ASSAY
6.23 AMYLASE ASSAY
6.24 PHENYL AMMONIA LYASE ASSAY

7. NUCLEIC ACIDS
7.1 ISOLATION OF GENOMIC DNA
7.2 SOLATION OF PLASMID
7.3 ISOLATION OF PLANT GENOMIC DNA
7.4 DETERMINATION OF DNA CONTENT
7.5 DETERMINATION OF CONCENTRATION AND PURITY OF THE DNA SAMPLE
7.6 ELECTROPHORESIS
7.7 RESTRICTION DIGESTION OF DNA
7.8 SOUTHERN BLOTTING
7.9 ISOLATION OF TOTAL RNA FROM PLANT TISSUE
7.10 ISOLATION OF RNA FROM YEAST
7.11 DETERMINATION OF RNA CONTENT (ORCINOL REAGENT METHOD)

8. MICROBIOLOGY
8.1 PREPARATION OF MEDIA
8.2 PREPARATION OF SLANTS, STABS AND POURING ON PETRIPLATES
8.3 GRAM STAINING OF BACTERIA
8.4 STAINING OF BACTERIAL SPORES
8.5 STREAK PLATE TECHNIQUE
8.6 ASSESSMENT OF AIR MICROFLORA
8.7 ENUMERATION OF BACTERIA IN SOIL
8.8 PREPARATION OF Escherichia coli COMPETENT CELLS AND ITS TRANSFORMATION
8.9 STUDY OF DIFFERENT GROWTH PHASES OF BACTERIAL POPULATION AND PLOT A BACTERIAL GROWTH CURVE
8.10 ANTIMICROBIAL SUSCEPTIBILITY TESTING (DISC DIFFUSION METHOD)
8.11 MINIMUM INHIBITORY CONCENTRATION (BROTH DILUTION METHOD)
8.12 MINIMUM BACTERICIDAL CONCENTRATIONS (MBC)
8.13 BIOCONVERSION OF TANNIC ACID TO GALLIC ACID BY Aspergillus niger
8.14 COMPARISON BETWEEN AEROBIC AND ANAEROBIC PROCESS ON ETHANOL PRODUCTION
8.15 DETERMINATION OF BIOMASS AND PACKED CELL VOLUME
8.16 CONJUGATION: Hfr MAPPING TO DETERMINE THE GENETIC DISTANCE BETWEEN GENES IN E. coli

9. DETECTION OF COMMON FOOD ADULTERANTS

10. APPENDIX
10.1 MOLARITY OF CONCENTRATED ACIDS
10.2 SOME TROUBLES AND REMEDIES FOR SDS-PAGE ANALYSIS
10.3 COMPOSITION OF BUFFERS AND MEDIA
10.4 TYPICAL IODINE NUMBERS
10.5 SAPONIFICATION VALUE OF COMMON FATS OR OIL

REFERENCES

PREFACE

A laboratory Text book of Biochemistry, Molecular Biology and Microbiology is intended to prepare the undergraduate, postgraduate and research students to perform basic experiments on various aspects of bioscience and biotechnology. Moreover, in the Semester system of teaching it is necessary to explore experiments which are not lengthy and easily completed within contact hours. Initially the book deals with dilutions, pH, buffers, units of measurements and calculations. This is followed by lab safety rules which is very important for any student working with chemicals for their and safety of others. This book emphasizes on principles, reagent preparations and procedures related to experiments, which will be handy for students from different scientific backgrounds. A number of methods are available in the literature for quantification of various molecules. This book does not present all the available methods but based on experience it contains commonly used methods, which students should know. The methods have been written in a manner for direct practical use in the laboratory.

This work has originated as a result of numerous requests from my students for eased out and explanatory methods pertaining to biochemistry, biotechnology, microbiology and others. The section on testing of adulterants is of much use for common mass because most of the food products we eat are adulterated. The approach is rather simple with the use of very easily available chemicals and the tests can be performed even in house. It is hoped that the reliable assays presented in this manual will help the students and research scholars to get to basics of experiments and various aspects associated with it.

Last but not the least I must thank my parents and especially my wife Neeti for constant persuasion and unending support which helped me out in completing this task effectively and efficiently. Suggestions for improvements are always welcome.

1. WATER, pH AND BUFFERS

WATER

Water is the major is the major component of living organisms. About 75-90% of a cell consists of water. The basis for much of these unusual properties is the presence of hydrogen bonds. The boiling point is 100oC, which is very high, compared to 161oC for methane having similar molecular weight as water. The density of water is highest at 4oC. This is unusual since most substances are denser in their solid state. Due to which the ice floats. Water has a high heat capacity, i.e. it will take a relatively large amount of energy to increase the temperature of water. As a result large bodies of water will experience far less severe temperature changes than the surrounding atmosphere. Water has high surface tension and considered as universal solvent.

pH

Because of the strong electronegativity of oxygen and H-bond formation the H-atom in water carries a partially positive charge. By chance the bonding pair in some water molecules may be shifted totally to the oxygen causing the ionzation of water, i.e. the formation of an OH- and H+ ion. However, the likely hood for this to happen is very small, to be exact: 1 mol of H2O contains 10-7mol of OH- and 10-7mol of H+. Water is said to dissociate and the degree of dissociation is a constant. It can be expressed as the dissociation constant (Kdiss).

illustration not visible in this excerpt

Since only 10-7mol of water (out of one mol water) are dissociated, the concentration of undissociated water [H2O] does not change significantly and can be taken into the dissociation constant Kdiss of water. The new constant KW is called the ion product of water and has a value of 10-14.

illustration not visible in this excerpt

Using the logarithm of the equation and multiplying by (-1) we obtain

illustration not visible in this excerpt

Using a convention to express: -log as p, the above equation can be rewritten as:

illustration not visible in this excerpt

In case of water pH = pOH = 7, i.e. water is neither acidic nor basic, it is neutral.

We measure the pH of a solution experimentally in two ways. In the first of these we use a chemical called an indicator, which is sensitive to pH. These substances have colours that change over a relatively short pH range (about two pH units) and can, when properly chosen, be used to determine roughly the pH of a solution. Two very common indicators are litmus, usually used on paper, and phenolphthalein, the most common indicator in acid-base titrations. Litmus changes from red to blue as the pH of a solution goes from about 6 to about 8. Phenolphthalein changes from colourless to red as the pH goes from 8 to 10. A given indicator is useful for determining pH only in the region in which it changes colour. Indicators are available for measurement of pH in all the important ranges of acidity and basicity. By matching the colour of a suitable indicator in a solution of known pH with that in an unknown solution, one can determine the pH of the unknown to within about 0.3 pH units.

The other method for finding pH is with a device called a pH- meter. In this device two electrodes, one of which is sensitive to [H+], are immersed in a solution. The potential between the two electrodes is related to the pH. The pH meter is designed so that the scale will directly furnish the pH of the solution. A pH- meter gives much more precise measurement of pH than does a typical indicator and is ordinarily used when an accurate determination of pH is needed. Some acids and bases undergo substantial ionization in water, and are called strong because of their essentially complete ionization in reasonably dilute solutions. Other acids and bases, because of incomplete ionization (often only about 1% in 0.1 M solution.), are called weak. Hydrochloric acid, HCl, and sodium hydroxide, NaOH, are typical examples of a strong acid and a strong base. Acetic acid, HC2H3O2, and ammonia, NH3, are classic examples of a weak acid and a weak base.

A weak acid will ionize according to the Law of Chemical Equilibrium:

illustration not visible in this excerpt

Ka is a constant characteristic of the acid HA. In solutions containing HA, the product of concentrations in the equation will remain constant at equilibrium independent of the manner in which the solution was made. A similar relation can be written for solutions of a weak base. The value of the ionization constant Ka for a weak acid can be found experimentally in several ways.

From the above equation, we can derive the Henderson-Hasselbach equation, commonly called the 'buffer equation', which relates the pH of solution to the pKa of the acid and the relative concentration of the undissociated acid and the conjugate base forms.

Solving first for the hydrogen ion concentration:

illustration not visible in this excerpt

Converting to the logarithmic form and multiplying by -1, we obtain:

illustration not visible in this excerpt

Defining operator 'p = -log' we have:

illustration not visible in this excerpt

Finally, by inverting the log ([HA]/[A-]) term, we obtain:

illustration not visible in this excerpt

This is the Henderson-Hasselbach equation. It is most useful in the preparation of buffers and in understanding how the concentration of the acid and conjugate base forms of a weak acid affect the pH.

When dissolved in water, salts of weak acids or weak bases furnish ions that tend to react to some extent with water, producing molecules of the weak acid or base and liberating some OH- or H+ ion to the solution.

BUFFERS

Buffers are solutions that contain mixtures of weak acids and bases that make them relatively resistant to pH change. Conceptually buffers provide a ready source of both acid and base to either provide additional H+ if a reaction (process) consumes H+, or combine with excess H+ if a reaction generates acid. The most common types of buffers are mixtures of weak acids and salts of their conjugate bases, for example, acetic acid/sodium acetate. In this system the dissociation of acetic acid can be written as

illustration not visible in this excerpt

Where, the acid dissociation constant is defined as Ka = [H+] [CH3COO-]/[H3COOH]. As discussed earlier rearranging and taking the negative logarithm gives the more familiar form of the Henderson-Hasselbach equation: Inspection of this equation provides several insights as to the functioning of a buffer. When the concentrations of acid and conjugate base are equal, log (1) = 0 and the pH of the resulting solution will be equal to the pKa of the acid. The ratio of the concentrations of acid and conjugate base can differ by a factor of 10 in either direction, and the resulting pH will only change by 1 unit. This is how a buffer maintains pH stability in the solution. To a first approximation, the pH of a buffer solution is independent of the absolute concentration of the buffer; the pH depends only on the ratio of the acid and conjugate base present. However, concentration of the buffer is important to buffer capacity.

1.1 PREPARATION OF ACETATE BUFFER

PRINCIPLE

Adding a 0.01 mL droplet of 1M HCl to 1L of pure water changes the pH of the water from 7 to about 5, which represents a 100 fold increase in [H+]. Such a huge change in pH would be intolerable to most biological systems, since even small changes in pH can dramatically affect the structures and functions of biological molecules. Regulation of the pH of the body fluids and tissues within limits consistent with life and normal function is provided by buffers. Buffers are defined as solutions that resist changes in the pH of a system upon addition of limited amounts of either acid or alkali. Buffers mixtures contain two substances, a conjugate acid and conjugate base. An "acidic" buffer contains a weak acid and its salt of a strong base. A "basic" buffer contains a weak base and its salt of strong acid. The action of buffers and their role in maintaining the pH of a solution can be explained with the aid of the Henderson–Hasselbach equation as discussed in the earlier section.

Together the two species (conjugate acid plus conjugate base) resist large changes in pH by partially absorbing addition of H+ or OH- ions to the system.

Buffered solutions do change in pH upon the addition of H+ or OH- ions. However, the change is much less than that which would occur if no buffer was present. The amount of change depends on the strength of the buffer and the conjugate base/conjugate acid ratio. The ability of a buffer to resist changes in pH is referred to as the "buffer capacity". Buffer capacity can be defined as the number of moles of H+ or OH- that must be added to one liter of the buffer in order to change the pH by one unit. The buffering capacity of the buffer solution is maximal when when the concentrations of conjugate acid and conjugate base are equal. In general, buffers should not be used at a pH greater or lower than 1 unit from their pKa.

In the laboratory, many biochemical reactions including those catalyzed by enzymes require pH control which is provided by buffers. The desired pH of the buffered solution determines which buffering compound is selected.

There are three practical methods to prepare a buffer:

A. If there is only one of the buffer components available, the buffer preparation is done by titration method.

a) In case of availability of weak acid or weak base only. An acidic buffer is made by titrating the weak acid by strong base, and a basic buffer is made by titrating the weak base by strong acid until the desired pH is obtained.
b) In case of availability of the salt of weak acid or base only. An acidic buffer is made by titrating the salt of weak acid by strong acid, and a basic buffer is made by titrating the salt of weak base by strong base until the desired pH is obtained.

Note: In this method, it can be calculating mathematically the volume of titrant required to prepare the desired buffer.

B. If both forms are available, using the buffer pKa, calculate the amounts (in moles) of acid/salt or base/salt present in the buffer at the desired pH. In calculation, convert the amount required from moles to grams using the molecular weight of that component, and then weight out the correct amounts. Or convert moles to volume if the stock is available in the liquid form.

C. Find a table of the correct amounts of acid/salt or base/salt required for different pH. Dissolve the components in slightly less water than is required for the find solution volume. Check that the pH and correct it if necessary. Add water to the final volume.

PROCEDURE

Acetate buffer can be prepared using acetic acid (HC2H3O2) (pKa = 4.756) and sodium acetate (NaC2H3O2), the sodium salt of the conjugate base. Note the relationship between the acid and its conjugate base in the equilibrium:

illustration not visible in this excerpt

First, make 100 mL of a buffer with pH = 5 using 5 mL of a 0.3M acetic acid solution. Using the Henderson-Hasselbach equation, the mass of sodium acetate needed to make the buffer is calculated as follows:

1. Calculate the concentration of acetic acid in the final 100 mL:

illustration not visible in this excerpt

2. Plug the values into HH and solve for unknown:

illustration not visible in this excerpt

3. Calculate the moles of A- needed to obtain this concentration in 100 mL:

illustration not visible in this excerpt

4. Calculate the mass of NaA needed to obtain this number of moles:

illustration not visible in this excerpt

Second, make 100 mL of a buffer also with pH = 5, but with a higher buffering capacity, using 5 mL of a 0.5 M acetic acid solution.

Although a buffer will resist a change in pH, eventually enough acid or base can be added to destroy it. The amount of acid or base needed to change the pH of a buffer is known as the "buffering capacity". Measure the buffering capacity of the buffers you make, monitoring the titration with a pH meter.

A. Preparation of pH = 5.0 Buffer A

1. Add about 50 mL of distilled water to a 100 mL beaker.
2. Use a volumetric pipette to add 5.0 mL of 0.30 M acetic acid to the beaker.
3. Accurately weigh a small beaker, add about 0.4 g sodium acetate (NaC2H3O2), and weigh the beaker again. Record these masses in your notebook to all the sig figs possible.
4. Insert a previously calibrated pH meter into the beaker.
5. Add a little bit of the sodium acetate at a time, swirling the beaker to dissolve it until the pH is 5.0. In theory, this will take 0.216 g of sodium acetate.
6. Weigh the beaker with the remaining sodium acetate.
7. Quantitatively transfer the buffer solution to a 100 mL volumetric flask.
8. Add distilled water up to the mark. Cap and invert the flask twice to mix.

B. Preparation of pH = 5.0 Buffer B

1. Calculate the mass of sodium acetate (NaC2H3O2) that must be added to make 100 mL of an acetic acid/acetate buffer at pH = 5.0, given that you will use 5.0 mL of 0.50 M acetic acid.

2. Make the buffer solution in a manner similar to that done in part A, starting with the mass of sodium acetate you calculated in B1 and using 0.50 M acetic acid. As you did in part A, weigh out a bit more sodium acetate than you will need.

1.2 TESTING OF BUFFER CAPACITY

1. Calibrate a pH meter carefully.
2. Using a volumetric pipette, transfer 25.0 mL of Buffer A (see previous experiment) into a 100 mL beaker or any other buffer.
3. Load a 50 mL burette with your Standardized NaOH solution.
4. Use the pH meter to monitor the titration of the buffer until the pH changes 1 unit. Run two titrations: a quick one and a careful one. For the quick one, add NaOH in 1 mL increments. For the careful one, use your best judgment in adding titrant.
5. Repeat Steps 2-4 for Buffer B (see previous experiment) or any other buffer.
6. Tabulate the data and analyse the result.

2. DILUTIONS, UNITS FOR MEASURING CONCENTRATION OF SOLUTIONS AND SPECTROPHOTOMETRIC CALCULATIONS

Solution is a homogeneous mixture of solute and solvent whose proportion varies within certain limits.

Solute is a substance which is present in small quantity in a solution.

Solvent is a substance which is present in large quantity in a solution.

Many solutions used in biochemistry are prepared by the dilution of a more concentrated stock solution. In preparing to make a dilution (or series of dilutions), you need to consider the goal of the procedure. This means that you need to consider both the desired final concentration and required volume of the diluted material. A simple equation allows the dilution to be calculated readily:

illustration not visible in this excerpt

Where, C1is the concentration of the initial solution; V1is the volume of the initial solution available to be used for dilution (this may not be the total volume of the initial solution, and instead may be a small fraction of the initial solution), C2is the desired final concentration, and V2is the desired final volume.

In most cases, the initial concentration and the final concentration are either known or are chosen in order to work correctly in the experiment being planned. The final volume is usually an amount that is chosen based on the amount required for a given experiment. This means that at least three of the required terms are either known or can be chosen by the experimenter.

For instance - You are setting up a standard curve and have a stock solution of 1000 μg/mL BSA, and for one of the points on the curve, you want 200 μl of 20 μg/mL. In this case, C1 = 1000 μg/mL; C2 = 20 μg/mL, and V2 = 200 μl. This leaves V1 as the unknown value (i.e. how much of the stock solution must be diluted to 200 μl final volume to yield the desired concentration). Rearranging the dilution equation gives:

illustration not visible in this excerpt

Thus, you need to dilute 4 μl of the stock solution to a final volume of 200 μl (i.e. by adding 196 μl). If, in the example, you wished to make a solution of 1 μg/mL, the same equation would indicate that you need 0.2 μl of the 1000 μg/mL stock solution for 200 μl of the final diluted sample. This is a problem: 0.2 μl is very difficult to measure accurately. You have two choices: change the final volume (i.e. if V2 is larger, then V1 must also increase), or perform serial dilutions (i.e. instead of diluting the stock solution by a factor of 1000 in one step, dilute the stock solution, and then make a further dilution of the diluted stock).

In many cases, while the final concentration is important, the final volume is not (as in the previous paragraph). In these cases, do what was explained in this example: use a convenient dilution: a dilution that involves volumes that are easily pipetted.

Pipetting 1.3333 μl is usually less accurate than pipetting 4 μl, both because 4 μl is a larger volume, and because it is difficult to set the pipet for 1.3333 μl. In this case, 4 μl is a convenient volume, while 1.3333 μl is not. In some cases, you may not know the actual starting concentration. If, for example, you need to measure the enzyme activity in a sample, and you find that the activity is too high to measure accurately, you will need to dilute the starting material. Since you don’t know the actual starting concentration, all you know is the concentration ratio between starting and final solutions. As long as you keep track of the concentration ratio in all of your dilutions, you can easily determine the enzyme activity in the initial solution, even though you cannot measure it directly.

Concentration ratios are frequently of considerable value. For example, you have a stock solution of buffer that contains 450 mM Tris-HCl, 10 mM EDTA, and 500 mM NaCl. You actually wish to use a final concentration of 45 mM Tris-HCl, 1 mM EDTA, and 50 mM NaCl. In each case the concentration of the final buffer is one-tenth that of the original. Simply performing a 1:10 dilution of the stock solution then gives the appropriate final concentration of each component. The stock solution of buffer is typically called a 10x stock, because it is ten-times more concentrated than the final, useful buffer.

The 1:10 dilution mentioned is performed by taking one part of the initial solution, and adding nine parts of solvent (usually water). This results in a final concentration that is ten-fold lower than the original.

SERIAL DILUTION

Now we know that dilution is the act of mixing a chemical with other substance, usually distilled water to make it lighter in composition. It is usually done if the chemical concentration is too high than the desired composition. Thus, serial dilution simply means a series of repeated dilution performed on the same chemical basically to change its concentration. After performing the dilution, we need to know how much difference are the diluted chemical and the initial, undiluted ones.

There are several benefits of performing serial dilution. Serial dilution comes in handy when the solution is too concentrated to be used in experiments or ingredients preparation. This is to ensure that the exact concentration can be obtained for the experiment to become success. Other than that, the diluted solution from this serial dilution can be used to count the concentration of the actual solution. By knowing the dilution factor of certain solution, the calculations of the concentration become easier and systematic.

This method is applicable in several fields, not only in chemistry. Serial dilution is a simple yet efficient technique to determine the number of cells or organisms in a concentrated sample. First, take a portion of the sample and does serial dilution on it. Repeat the steps until the cells can be observed under the microscope when the diluted sample was observed. Then, count the dilution factor and times it with the actual volume of the sample. Other than that, medicine administration require suitable dose for each patient with variable needs. This is where serial dilution is useful.

Some of the terminologies are related to the serial dilution are as follows:

An aliquot/sample is a measured sub-volume of original sample.

Diluent is material with which the sample is diluted

Dilution factor (DF): ratio of final volume/aliquot volume (final volume = aliquot + diluent)

Dilution = aliquot volume/total volume = aliquot volume/aliquot volume + diluent volume

For example: What is the dilution factor if you add 0.1 mL aliquot of a specimen to 9.9 mL of diluent?

1. The final volume is equal to the aliquot volume plus the diluent volume: 0.1 mL + 9.9 mL = 10 mL

2. The dilution factor is equal to the final volume divided by the aliquot volume: 10 mL/0.1 mL = 1:100 dilution (10 2)

The dilution for this problem = aliquot volume/final volume = 0.1/(0.1 + 9.9) = 0.01 or 10 -2.

Additional examples on dilution and serial dilutions

A. One mL of a sample was mixed with 99 mL of buffer. One mL of this was plated (using the pour plate method) in nutrient agar. After incubation, 241 colonies were present on the plate. How many colony-forming units were present per mL of the original sample? State your answer in CFU/mL.

The dilution used was [Abbildung in dieser Leseprobe nicht enthalten]. One mL of the dilution contained 241 colony-forming units. How much did one mL of the original sample contain? Obviously more than 241 colony-forming units! To arrive at the correct number, either divide the colony-forming units/mL of the dilution by the dilution (241 colony-forming units/mL divided by [Abbildung in dieser Leseprobe nicht enthalten] of original sample) or multiply by the dilution factor. The dilution factor is defined as the inverse of the dilution, therefore in this example the dilution factor would be 1/10-2 = 102. Multiplying by the dilution factor: 241 colony-forming units[Abbildung in dieser Leseprobe nicht enthalten] of original sample. Either way, the answer you get is the same.

B. One mL of a sample was mixed with 99 mL buffer. One-tenth of amL of this was plated on nutrient agar. After incubation, 142 colonies were present on the plate. How many colony-forming units were present per mL of the original sample?

In this case, the mathematics of the dilutions is the same, but the number of colonies counted on the plate (the number of colony-forming units in only 0.1 mL of the dilution).

Remember, you should always report your answers as CFU/mL. Therefore, the number of colony-forming units per mL of dilution is 142 colonies divided by 0.1 mL plated, or 142 x 10 = 1420 colony-forming units/mL of dilution. Dividing this number by the 10-2 dilution results in a final answer of 1420 x 102 or 1.42 x 105 colony-forming units/mL of original sample.

C. An overnight culture of Escherichia coli is used as a sample. One mL of this culture is added to a bottle containing 99mL of buffer. This dilution is mixed well (as all dilutions are!), and one mL of this is mixed in 9 mL of buffer. This second dilution is diluted by three successive 1/10 dilutions. The last (fifth) dilution is plated, i.e. 0.1 mL is plated on nutrient agar. After incubating the plate, 56 colonies are counted. How many colony-forming units were present per mL of E. coli culture?

To solve this problem, try writing out the procedure. Then multiply the successive dilutions together (here is where scientific notation comes in handy, since it is easy to add the exponents without 'losing a zero').

illustration not visible in this excerpt

OR

illustration not visible in this excerpt

The number of colony-forming units per mL of the dilution can then be divided by the final dilution:

illustration not visible in this excerpt

Note that the 0.1 mL that was plated is treated as another 1/10 dilution. Some microbiologists treat it as such in their computations rather than dividing the colony count by the number of milliliters plated, as done here. Whichever way it is done, the answer will be the same.

Units for measuring concentration of solutions

1. Percentage by mass
2. Normality
3. Molarity
4. Molality
5. Mole-fraction

1. Percentage by mass

It is defined as the number of grams of solute present in 100grams of solution.

For example: A 20% solution of NaOH by weight contains 20 parts by mass of NaOH dissolved in 80 parts by mass of water. Generally this unit is used to prepare solutions of approximate concentration.

2. Normality (N)

Normality is the number of gram equivalents of solute present in1000 mL of solution. Normality is represented by the symbol 'N'

Decinormal Solution (Normality = 0.1N)

A decinormal solution contains one tenth of a gram equivalent of solute in one litre of solution.

N=Mass of Solute x 1000/Equivalent mass of solute x Volume of solution

For example:

Calculate the normality of an oxalic acid solution containing 3.2g of oxalic acid in 2 litres of solution.

illustration not visible in this excerpt

Mole or gram mole is the molecular mass of a substance expressed in grams.

Number of moles = Mass of Solute in grams/Molecular mass of solute

3. Molarity (M)

Molarity is the number of moles of solute present in 1000 mL or one litre of solution.

Molarity is represented by the symbol 'M'.

Molar solution (Molarity =1M)

A Molar solution contains one mole of solute in one litre of solution.

M = Mass of Solute x 1000/Molecular mass of solute x Volume of solution

We can also have another formula which relates Molarity and Normality

Molarity x Molecular mass = Normality x Equivalent mass.

For example:

Calculate the molarity of a solution containing 50g of sugar (C12H22O11) in 400 mL of solution.

illustration not visible in this excerpt

Another example:

Find the normality of a solution of sulfuric acid whose molarity is equal to 5M.

illustration not visible in this excerpt

4. Molality

Molality is the number of moles of solute present in 1000g (or) 1Kgof the solvent. Molality is represented by the symbol 'm'.

A molal solution contains one mole of solute in one Kg of solvent.

Molality = Mass Solute x 1000/Molecular mass of solute x Mass of solvent

For example:

Calculate the molality of a solution containing 15g of methanol (CH3OH) in 300g of solvent.

illustration not visible in this excerpt

5. Mole-fraction (X)

Mole-fraction of solvent (X1)

Mole-fraction of the solvent is the ratio between the number of moles of solvent (n1) and the total number of moles of solute and solvent (n1+ n2) present in solution.

Mole-fraction of solvent = Number of Moles of solvent/Total number of Moles present in solution

illustration not visible in this excerpt

Mole-fraction of the solute is the ratio between the number of molesof solute (n2) and the total number of moles of solute and solvent (n2+n1) present in solution.

Mole-fraction of solute= Number of Moles of solute/ Total number of Moles present in solution

illustration not visible in this excerpt

Important to note:

Number of moles = Mass in gram/ Molecular Mass

In any solut1ion, the sum of mole-fraction of solute and solvent is equal to one

illustration not visible in this excerpt

For example:

Find the number of moles of solute and solvent in a solution containing 9.2g of ethyl alcohol in 180 g of water.

illustration not visible in this excerpt

Number of moles of ethyl alcohol n = Mass / Molecular mass

illustration not visible in this excerpt

Mole-fraction of solvent = Number of Moles of solvent/ Total number of Moles present in solution

illustration not visible in this excerpt

Mole fraction of solute = Number of Moles of solute/ Total number of Moles present in solution

illustration not visible in this excerpt

SPECTROPHOTOMETRIC CALCULATION

First of all prepare range of standard solutions (mg/mL or µg/mL) which are processed according to the protocol and draw a standard graph. Then take the absorbance of the sample. Extrapolate it from the standard graph and calculate the amount according to the following equation.

e.g. While determining the flavonoid content the equation used is:

Flavonoid content % = Quercetin equivalent [Abbildung in dieser Leseprobe nicht enthalten]

Quercetin (standard flavonoid compound) equivalent to be checked from standard graph of Quercetin.

illustration not visible in this excerpt

The above equation can be followed in all the forthcoming spectrophotometric assays.

3. LAB SAFETY RULES

1. Conduct yourself in a responsible manner at all times in the laboratory.
2. Follow all written and verbal instructions carefully. If you do not understand a direction or part of a procedure, ask your teacher before proceeding with the activity.
3. Never work alone in the laboratory. No student may work in the science classroom without the presence of the teacher.
4. When first entering a science room, do not touch any equipment, chemicals, or other materials in the laboratory area until you are instructed to do so.
5. Perform only those experiments authorized by your teacher. Carefully follow all instructions, both written and oral. Unauthorized experiments are not allowed.
6. Do not eat food, drink beverages, or chew gum in the laboratory. Do not use laboratory glassware as containers for food or beverages.
7. Be prepared for your work in the laboratory. Read all procedures thoroughly before entering the laboratory. Never fool around in the laboratory. Horseplay, practical jokes, and pranks are dangerous and prohibited.
8. Always work in a well-ventilated area.
9. Observe good housekeeping practices. Work areas should be kept clean and tidy at all times.
10. Be alert and proceed with caution at all times in the laboratory. Notify the teacher immediately of any unsafe conditions you observe.
11. Dispose of all chemical waste properly. Never mix chemicals in sink drains. Sinks are to be used only for water. Check with your teacher for disposal of chemicals and solutions.
12. Labels and equipment instructions must be read carefully before use. Set up and use the equipment as directed by your teacher.
13. Keep hands away from face, eyes, mouth, and body while using chemicals or lab equipment. Wash your hands with soap and water after performing all experiments.
14. Experiments must be personally monitored at all times. Do not wander around the room, distract other students, startle other students or interfere with the laboratory experiments of others.
15. Know the locations and operating procedures of all safety equipment including: first aid kit(s), and fire extinguisher. Know where the fire alarm and the exits are located.
16. Know what to do if there is a fire drill during a laboratory period; containers must be closed, and any electrical equipment turned off.
17. Any time chemicals, heat, or glassware are used, students will wear safety goggles. NO EXCEPTIONS TO THIS RULE!
18. Contact lenses may be not be worn in the laboratory.
19. Dress properly during a laboratory activity. Long hair, dangling jewellery, and loose or baggy clothing are a hazard in the laboratory. Long hair must be tied back, and dangling jewellery and baggy clothing must be secured. Shoes must completely cover the foot. No sandals allowed on lab days.
20. A lab coat or smock should be worn during laboratory experiments.
21. Report any accident (spill, breakage, etc.) or injury (cut, burn, etc.) to the teacher immediately, no matter how trivial it seems. Do not panic.
22. If you or your lab partner is hurt, immediately (and loudly) yell out the teacher's name to get the teacher's attention. Do not panic.
23. If a chemical should splash in your eye(s) or on your skin, immediately flush with running water for at least 20 minutes. Immediately (and loudly) yell out the teacher's name to get the teacher's attention.
24. All chemicals in the laboratory are to be considered dangerous. Avoid handling chemicals with fingers. Always use a tweezer. When making an observation, keep at least 1 foot away from the specimen. Do not taste, or smell any chemicals.

25. Check the label on all chemical bottles twice before removing any of the contents. Take only as much chemical as you need.
26. Never return unused chemicals to their original container.
27. Never remove chemicals or other materials from the laboratory area.
28. Never handle broken glass with your bare hands. Use a brush and dustpan to clean up broken glass. Place broken glass in the designated glass disposal container.
29. Examine glassware before each use. Never use chipped, cracked, or dirty glassware.
30. If you do not understand how to use a piece of equipment,
ASK THE TEACHER FOR HELP!
31. Do not immerse hot glassware in cold water. The glassware may shatter.
32. Do not operate a hot plate by yourself. Take care that hair, clothing, and hands are a safe distance from the hot plate at all times. Use of hot plate is only allowed in the presence of the teacher.
33. Heated glassware remains very hot for a long time. They should be set aside in a designated place to cool, and picked up with caution. Use tongs or heat protective gloves if necessary.
34. Never look into a container that is being heated.
35. Do not place hot apparatus directly on the laboratory desk. Always use an insulated pad. Allow plenty of time for hot apparatus to cool before touching it.

Fig. 3.1: Laboratory Safety Symbols

(http://www.abss.k12.nc.us/cms/lib02/NC01001905/Centricity/Domain/1835/Laboratory%20Safety%20Symbols%20and%20Rules.docx)

illustration not visible in this excerpt

4. CARBOHYDRATES

THEORY

Along with proteins and fats, carbohydrates, or sugars, are one of the three main classes of food used to sustain life. They are essential components of all living organisms and constitute the most abundant class of biological molecules. The classification carbohydrate stems from the general molecular formula for monosaccharides, (C·H2O)n where n ≥ 3, which implies that these compounds are hydrates of carbon. However, carbohydrates are not true hydrates in the chemical sense. Carbohydrates, chemically, are polyhydroxy aldehydes (-CHO) or ketones (C=O) or compounds which upon hydrolysis yield these compounds. Note that each carbon in a polyhydroxy aldehyde or ketone structure, except for the carbonyl functional group, bears a hydroxyl (-OH) functional group. Polyhydroxy aldehydes and ketones with the same number of carbons are structural isomers of each other.

Carbohydrates mat be classified as monosaccharide, disaccharide and polysaccharides. Monosaccharides are further classified based upon the carbonyl functional group present and the total number of carbon atoms in the structure. The polyhydroxy aldehyde on the previous page contains five carbons and is classified as an aldopentose, where the prefixes aldo- and pent- indicate the aldehyde functional group and five carbon atoms, respectively. The polyhydroxy ketone structure is a ketopentose, where keto-indicates a ketone functional group. Disaccharides are carbohydrates which upon hydrolysis (reaction with water) yield two monosaccharide structures. The most abundant disaccharide is sucrose, which is hydrolyzed to D-glucose and D-fructose. The two sugar units are covalently joined through an oxygen atom at carbons 1 and 2 on glucose and fructose, respectively e.g. maltose and lactose. The covalent link between the two sugar units is referred to as an (α1→β2) glycosidic linkage. The glycosidic linkages are destroyed when disaccharides undergo hydrolysis back to the individual monosaccharides. Polysaccharides consist of large numbers of monosaccharides joined by glycosidic linkages. When the monosaccharide units are all identical, the molecule is referred to as a homopolysaccharide. If the monosaccharide units differ, the molecule is classified a heteropolysaccharide. Starch is a homopolysaccharide synthesized by plants for the storage of α-D-glucose units and serves as a source of carbohydrates in animal diets. Plant starch occurs as a mixture of two different polymeric structural units. α- amylose is a linear polymer of several thousand glucose units joined by α(1→4) glycosidic bonds. Amylopectin is a branched polymer with glucose connected in linear chains by α(1→4) bonds and by α (1→6) bonds at the branch points which occur on the average every 24 to 30 glucose units. Containing up to a million glucose residues, amylopectin is the larger of the two polymeric structures. Glycogen, the storage polysaccharide for glucose in animals and humans is structurally similar to amylopectin except that branching occurs every 8 to 12 glucose residues.

4.1 QUALITATIVE TESTS FOR CAROHYDRATES

The chemical reactions of carbohydrates are largely that of the hydroxy and carbonyl groups. The aldehyde and α-hydroxy ketone (a ketone that has an -OH on a carbon next to the carbonyl group) structural units in sugars undergo mild oxidation to carboxylic acids. In solution, the ring structures of sugars are able to open at the anomeric carbons to form the open-chain aldehyde or α -hydroxy ketone form, which undergoes oxidation in the presence of a mild oxidizing agent. Such sugars are referred to as reducing sugars. However, when the anomeric ring carbon is involved in a glycosidic linkage, the ring is locked in place and cannot open for oxidation to occur. All monosaccharides are reducing sugars since no glycosidic linkages exist in their structures. The disaccharides maltose and lactose are reducing sugars since only one of the C1 anomeric carbons is involved in a glycosidic linkage, while sucrose is a nonreducing sugar since both the C1 and C2 anomeric carbons participate in the glycosidic linkage. Starch and glycogen are also nonreducing sugars because of their size.

REQUIREMENTS

(I) Fehling’s reagent A: Dissolve 34.65 g copper sulphate in distilled water and make up to 500 mL.
(II) Fehling’s reagent B: Dissolve 125 g potassium hydroxide and 173 g Rochelle salt (Potassium sodium tartrate) in distilled water and make up to 500 mL.
(III) Iodine solution: Add a few crystals of iodine to 2% potassium iodide solution till the colour becomes deep yellow.
(IV) Benedict’s qualitative reagent: Dissolve 173 g sodium citrate and 100 g sodium carbonate in about 500 mL water. Heat to dissolve the salts and filter, if necessary. Dissolve 17.3 g copper sulphate in about 100 mL water and add it to the above solution with stirring and make up the volume to 1 L with water.
(V) Seliwanoff’s reagent: Dissolve 0.05 g resorcinol in 100 mL dilute (1:2) hydrochloric acid.
(VI) Barfoed’s reagent: Dissolve 24 g copper acetate in 450 mL boiling water. Immediately add 25 mL of 8.5% lactic acid to the hot solution. Mix well, Cool and dilute to 500 mL.
(VII) Bial’s reagent:Dissolve 1.5 g orcinol in 500 mL of concentrated HCl and add 20 to 30 drops of 10% ferric chloride.
(VIII) Test tubes, carbohydrate sample, water bath, test tube stand

A. MOLISH TEST

PRINCIPLE

The Molisch test is used to determine the presence of carbohydrates, regardless of structure. The Molisch reagent provides the condensation reagent, α-naphthol in alcohol. Concentrated sulfuric acid acts as the dehydrating reagent. It also serves as a catalyst for the hydrolysis of di- and polysaccharides. The presence of a purple ring at the interface between the carbohydrate solution and concentrated sulfuric acid is a positive indication of a carbohydrate structure. Monosaccharides react more quickly than di- and polysaccharides, which must slowly undergo hydrolysis before they react with the Molisch reagent.

PROCEDURE

1. Add two drops of Molisch’s reagent to about 2 mL of test solution and mix well.
2. Incline the tube at about 45° angle and add about 1 mL of concentrated sulfuric acid along the sides of the tube.
3. Place in a rack and note the time for the appearance of a red to purple ring at the interface of the two layers. Aldopentoses, ketopentoses, and ketohexoses react more rapidly than aldohexoses and disaccharides.

B.BENEDICT’S TEST

PRINCIPLE

Benedict's reagent reacts with reducing sugars to form inorganic precipitates which are readily detected by visual observation. In the reaction between a reducing sugar and Benedict's reagent, copper (II) ion is reduced to copper(I) by the aldehyde functional group. The formation of a red to orange precipitate indicates the presence of a reducing sugar.

PROCEDURE

1. Place 5 drops of sample solution in a test tube.
2. Add 20 drops of Benedict's reagent to each solution and mix thoroughly.
3. Place test tube in boiling water for 1-2 min and note the time for the appearance of a red to reddish orange precipitate.

C. TOLLENS'S TEST

PRINCIPLE

The Tollens's test also involves a mild oxidizing agent. In the Tollens's test, silver (I) ion is reduced to metallic silver. The formation of a silver mirror on the inside of the test tube indicates the presence of a reducing sugar. The Tollens's reagent must be freshly prepared to produce the proper results.

PROCEDURE

1. Place 20 drops of 5% silver nitrate into a 100 mL beaker and add 1 drop of 3 M NaOH to the beaker.
2. Continue to add 2% ammonia with stirring to the beaker until the brownish precipitate of silver oxide just disappears.
3. Take this solution and add 5 drops of each sugar solution. Do not mix.
4. Place the test tubes in the boiling water bath for 2 minutes and observe formation of a silver mirror on the inside of the test tube.

D. BARFOED’S TEST

PRINCIPLE

Barfoed's reagent also uses the reduction of copper (II) ion to red Cu2O as an indication of a reducing sugar. However, it not as reactive as Benedict's reagent, and the rate at which the red precipitate forms can be used to distinguish between mono-saccharides and disaccharides. The appearance of red Cu2O within 2-3 minutes is a positive test for a monosaccharide; disaccharides produce the precipitate in approximately 10 minutes.

PROCEDURE

1. To 1 mL of the test solution add about 2 mL of Barfoed’s reagent.
2. Boil it for one minute and allow it to stand for a few minutes.
3. Observe the formation of brick red colour.

E. BILAL’S TEST

PRINCIPLE

The ability of the hydroxy groups to undergo dehydration reactions allows for identification of certain aldoses and ketoses. Concentrated acid is used to dehydrate the sugars, which produce colourless dehydration products. These dehydration products then react with a condensation reagent to form coloured condensation products, which offer a visual means of positive identification. Bial's test is used to distinguish a pentose from a hexose structure. The dehydrating agent is concentrated HCl, and the condensation reagent is orcinol in the presence of iron (III) chloride. Pentoses undergo dehydration to furfural, a cyclic aldehyde, which further reacts with orcinol to give a blue-coloured condensation product. All other coloured products are negativeindicators for the presence of a pentose.

PROCEDURE

1. Place 5 drops of test solution in labeled test tube.
2. Add 20 drops of Bial's reagent to each test tube. Mix thoroughly.
3. Place test tubes in boiling water and remove after 1 minute.
4. A dark blue colour indicates a pentose.

F. SELIWANOFF TEST

PRINCIPLE

The Seliwanoff test differentiates between ketohexoses and aldohexoses. It also uses concentrated HCl for the dehydrating agent, but the condensation agent is resorcinol. Ketohexoses are dehydrated to 5-hydroxymethylfurfural, which undergoes condensation with resorcinol within two minutes to form a red-coloured product. The appearance of a peach colour is not a positive test for a ketohexose.

PROCEDURE

1. Place 5 drops of test solution in labeled test tube.
2. Add 20 drops of Seliwanoff reagent and mix thoroughly.
3. Place the test tubes in boiling water and note the time for the appearance of a cherry red colour. Poly- and disaccharides hydrolyze slowly in the presence of acid and should take more time to give a positive test.
4. After two minutes remove from the water and place in a rack and record the results.

G. MUCIC ACID TEST

PRINCIPLE

This test is one in which concentrated HNO3 is heated along with an aldose sugar to give a dicarboxylic acid. Nitric acid is able to oxidize the terminal groups of aldoses, but leaves the secondary hydroxyl groups unchanged. The dicarboxylic acid formed from galactose is called mucic acid and is insoluble in cold aqueous solution. Those acids formed from the other common sugars are soluble in H20. Thus the formation of the insoluble precipitate is an indication of the presence of galactose. Lactose will also yield a mucic acid, due to hydrolysis of the glycosidic linkage between its glucose and galactose subunits.

PROCEDURE

1. Add 1 mL of concentrated nitric acid to 5 mL of the solution to be tested and mix well.
2. Heat on a boiling water bath until the volume of the solution is reduced to about 1 mL.
3. Remove the mixture from the water bath and let it cool at room temperature overnight.
4. The presence of insoluble crystals in the bottom of the tube indicates the presence of mucic acid.

H. IODINE TEST

PRINCIPLE

Iodine (iodine-potassium iodide, I2KI) staining distinguishes starch (a polysaccharide) from monosaccharides, disaccharides, and other polysaccharides. The basis for this test is that starch is a coiled polymer of glucose. Iodine interacts with these coiled molecules and becomes bluish black. Other non-coiled carbohydrates do not react with iodine. Therefore, a bluish black colour is a positive test for starch, and a yellowish brown colour (i.e., no colour change) is a negative test for starch. Glycogen, the common polysaccharide in animals, has a slight difference in structure and produces only an intermediate colour reaction. Test each of the known sugars for the presence of starch.

PROCEDURE

1. Add few drops of Iodine-KI reagent to 1 mL of test sample.
2. Appearance of bluish black colour indicates the presence of starch.
3. If starch is not present, then the colour will stay orange or yellow.

I. FEHLING’S TEST

PRINCIPLE

The blue alkaline cupric hydroxide present in Fehling’s solution, when heated in the presence of reducing sugars, gets reduced to yellow or red cuprous oxide and it gets precipitated. Hence, formation of the coloured precipitate indicates the presence of reducing sugars in the test solution.

PROCEDURE

1. To 1 mL of Fehling’s solution ‘A’ add 1 mL of Fehling’s solution ‘B’ and a few drops of the test solution.
2. Boil for a few minutes.
3. Observe the change in colour.

Tips to Care

1. Wear suitable protective clothing, gloves, and eye/face protection.
2. BENEDICT'S REAGENT: It is eye, skin, and respiratory irritant harmful if swallowed; may cause gastrointestinal discomfort.
3. BARFOED'S REAGENT: It is harmful if swallowed, may cause gastrointestinal discomfort, corrosive to skin and eyes.
4. BIAL'S REAGENT: It is corrosive to skin and eyes; harmful if ingested.
5. SELIWANOFF REAGENT: It is corrosive to skin and eyes, harmful if ingested.
6. AMMONIUM HYDROXIDE: contact may cause tissue damage to skin, eyes, mucous membrane, gastrointestinal & respiratory mucosa.
7. SILVER NITRATE: Damage to the eyes and can burn skin, inhalation can irritate respiratory passages and mucous membranes, ingestion can cause severe abdominal pain and gastroenteritis, that may be fatal.

4.2 DETERMINATION OF REDUCING SUGARS CONTENT (NELSON-SOMOGYI METHOD)

PRINCIPLE

Carbohydrates are the important components of storage and structural materials in the plants. They exist as free sugars and polysaccharides. Sugars having free aldehyde or keto group are called reducing sugars e.g. glucose, galactose, lactose and maltose. The Nelson-Somogyi method is one of the most commonly used methods for the quantification of reducing sugars. Reducing sugars when heated with alkaline copper tartrate reduce the copper from the cupric to cuprous state and thus cuprous oxide is formed. Cuprous oxide when reacts with arsenomolybdic acid there is reduction of molybdic acid to molybdenum giving blue colour. The blue colour is measured using spectrophometer at 620 nm.

REQUIREMENTS

(I) Copper Tartarate Solution (A): Anhydrous sodium carbonate (2.5 g), sodium bicarbonate (2 g), potassium sodium tartrate (2.5 g) and anhydrous sodium sulphate (20 g) were mixed in 80 mL water and diluted to 100 mL.
(II) Copper Sulphate Solution (B): Dissolve 15 g of copper sulphate in 100 mL of distilled water and add one or two drops of sulfuric acid. Mix 96 mL of (A) and 4 mL of (B) before use.
(III) Arsenomolybdate reagent: Dissolve 2.5 g ammonium molybdate in 45 mL distilled water. Add 2.1 mL sulfuric acid and mix well and add 0.3 g disodium hydrogen arsenate dissolved in 25 mL water. Mix well and incubate at 37°C for 48 hours and stored in an amber coloured bottle.
(IV) Standard stock solution of Glucose:50 mg in 50 mL distilled water. Working solution: Dilute 10 mL of stock to 100 mL with distilled water (0.1 mg/mL).

PROCEDURE

1. Weigh 100 mg of the sample and extract sugars with hot 80% ethanol at least twice (5 mL each).
2. Collect the supernatant and evaporate on a water bath at 80°C.
3. Add 10 mL distilled water to dissolve sugars completely.
4. Pipette out aliquots (0.1/ 0.2 mL) to separate test tubes.
5. Pipette out 0.2, 0.4, 0.6, 0.8 and 1 mL of the working standard solution into a series of test tubes.
6. Make up the volume in sample tube and standard tubes to 2 mL with distilled water.
7. Blank was prepared using 2 mL distilled water.
8. Add 1 mL of alkaline copper tartrate reagent to each tube.
9. Incubate the tubes in boiling water for 10 minutes.
10. Let the tubes to cool and add 1 mL of arsenomolybolic acid reagent to all the tubes.
11. Make up the volume to 10 mL with distilled water in each tube.
12. Blue coloured solution was read at 620 nm after 10 min.
13. Calculate the amount of reducing sugars present in the sample from standard graph.

[...]