BO101 Lab Manual 2024-25 PDF

1st Year Biology Practical Manual Semester 1: Biochemistry Ollscoil na Gaillimhe University of Galway Student Name: Student Number:  General laboratory information Attendance The practical classes are an essential part of the BO101 Biology course. A pre-...

1st Year Biology Practical Manual Semester 1: Biochemistry Ollscoil na Gaillimhe University of Galway Student Name: Student Number:  General laboratory information Attendance The practical classes are an essential part of the BO101 Biology course. A pre- practical talk is given in the teaching laboratory before each practical. Attendance at the pre-practical talks and the laboratory sessions is compulsory. Attendance is recorded at each practical class. If you are ill, you must email a medical certificate to the [email protected] Presentation of Work The practical manual must be read thoroughly before the practical. All data observations must be recorded directly into this manual. There is sufficient time in the practical class to record all data and complete all calculations. The completed workbook must be shown to the demonstrator before leaving the laboratory. Assessment Your practical classes will be examined as part of the Biology BO101 end of semester exam.. Laboratory Rules: 1. Laboratory coats (buttoned-up) must be worn at all times in the laboratory. 2. Some of the chemicals that you will be using are potentially poisonous. For this reason eating and drinking in the laboratory are forbidden. 3. Wash your hands frequently and always before leaving the laboratory. 4. Do not put any solids (e.g. paper, pipette tips) into the sinks. Waste bins are provided for this purpose. 5. If it is necessary to pour acid or alkali into the sink, let the water run freely to dilute the acid or alkali and wash it away. 6. Many of the reagents used are shared with your classmates. Both common sense and courtesy are required in their use. 7. If you are required to collect reagents, bring suitably sized clean dry labelled containers (flasks, beakers and test tubes) to the reagent bench and take only a sufficient amount of reagent for your experiment. In most cases, a reasonable amount will be indicated on or near the reagent bottle. 8. Test tubes should be labelled with an indelible marker, which you must provide yourself, in order to avoid confusion when dealing with colourless solutions. 9. ‘Parafilm’ should be used to seal test tubes before mixing by inversion to prevent spillage of solutions. 10. No pipetting of solutions by mouth is permitted. Use the Pasteur pipettes or autopipettes provided. 11. At the completion of each day’s laboratory work, all glassware must be returned to collection boxes for dirty glassware. 12. The wearing of safety glasses is recommended. These may be purchased from the ‘Student Union Shop’. 13. Personal belongings such as coats and bags should not be brought into the teaching laboratory but stored in your locker. 14. All personal accidents, however trivial, should be reported immediately to the class supervisor. Fire Safety: The following precautions are most important: 1. Locate laboratory fire extinguishers, fire blanket and first aid kit. Remember their location. 2. If inflammable volatile solvents are spilled, clean them up immediately. Dispose of such liquids by flushing them down a sink, in a fume hood, with plenty of water. 3. Small fires involving organic solvents, especially on persons or their clothing, can often be put out by smothering with a fire blanket or item of clothing to hand e.g. a coat. 4. For all other fires use a CO2 extinguisher or an asbestos blanket. 5. Do not use water on an organic solvent fire. It usually serves spread the fire. 6. Note the location of all fire exits to the laboratory. 7. In the case of fire, leave the building quickly and calmly by the nearest fire exit and assemble at the designated ‘Fire Assembly Point’. Poisons and Corrosive Liquids: while solutions used in first year labs are generally non-toxic, all solutions should be treated as potentially poisonous. 1. Some of the chemicals used are potentially poisonous even in small quantities, while others are potentially corrosive and liable to damage skin, clothing and footwear. 2. Handle all solutions with care. 3. Safety glasses should be worn to protect your eyes when handling corrosive chemicals. 4. If a corrosive substance splashes on your skin e.g. acid, rinse well with running water immediately and report to the class supervisor. Boiling Water Baths 1. Care should be taken when working with boiling water baths. 2. Enlist the help of a demonstrator for placing or removing test tubes in the water bath. Glassware 1. Care should be taken when handling glassware especially glass pipettes. 2. Broken glass should not be touched by hand. Use a brush and pan. 3. Broken glass should not be thrown in waste bins but discarded carefully in the ‘Sharps’ container provided. 3  1st Year Biochemistry Practical Table of Contents Page Introduction 5 Working with Units and Numbers 6 Lab. 1: Making Solutions, pH and Autopipettes 11 Lab. 2: Proteins and Nucleic acids 21 Lab. 3: Lipids and Carbohydrates 30 TIMETABLE (Provisional) Dates Lab 1: Making Solutions, pH and Autopipettes Oct. 2nd, 3rd, 4th Lab 2: Proteins and Nucleic Acids Oct. 9th, 10th, 11th Lab 3: Lipids and Carbohydrates Oct. 16th, 17th, 18th 4 Introduction Biochemistry is the science concerned with the chemical basis of life. It is the study of the various molecules that occur in living cells and their chemical reactions. An understanding of biochemistry is important in science and medicine. A fundamental understanding of the molecules that make up cells is fundamental to the understanding of how all cells, including bacteria, plant, animal and human cells, function. Since all diseases have a biochemical basis, biochemical tests are used extensively in medicine for diagnosis, prognosis, monitoring and screening of disease. This course in practical biochemistry is designed to introduce you to the four major types of biomolecule that are found in living cells: lipids, carbohydrates, proteins and nucleic acids. It will serve to complement and supplement your theoretical course. Its success depends largely on your approach to the laboratory exercises. It is essential that you obtain a clear rationale of the experiment through reading the relevant sections in textbooks and the notes for each practical. You are expected to have a complete knowledge of every phase of the experiment including the work carried out by your partner. Big Ideas  There are four classes of biological molecules: Proteins, lipids, carbohydrates and nucleic acids.  All biological molecules are, for the most part, made up of a small number of elements: Carbon, oxygen, nitrogen, phosphorus and sulphur. Other elements are found in smaller amounts.  Simple tests can be used to detect the presence of proteins, lipids, carbohydrates and nucleic acids in given samples. Learning goals/objectives for students: Students will be able to:  name the four biological macromolecules and their building blocks.  test samples for the presence of lipids, proteins, and simple and complex sugars.  understand how to isolate DNA. Working with Units and Numbers Learning Outcomes include an understanding of:  units and the ability to convert between units  scientific notation Introduction One fundamental skill required in Biochemistry is the ability to use units and to be able to convert one unit to another. The units used in modern science are those of 5 the metric system. Prior to the introduction of the metric system, dozens of different units were commonly used throughout the world e.g. length could be measured in feet, inches, miles, furlongs etc. The lack of common standards led to much confusion. At the end of the 18th century, the French government sought to alleviate the problem by devising a system of measurement that could be used worldwide. In 1790, the French National Assembly commissioned the French Academy of Science to design a simple decimal-based system of units which came to be known as the metric system. The International System of Units (SI) specifies a set of unit prefixes known as SI prefixes or metric prefixes. An SI prefix is a name that precedes a basic unit of measure to indicate a decimal multiple or fraction of the unit. Each prefix has a unique symbol that is prepended to the unit symbol. The SI prefixes are standardized by the International Bureau of Weights and Measures from 1960 to 1991. SI prefixes are used to reduce the number of zeros shown in numerical quantities before or after a decimal point. For example, an electrical current of 0.000000001 ampere, or one- billionth (short scale) of an ampere, is written by using the SI-prefix nano as 1 nanoampere or 1 nA. In 1960, the metric system was officially named the Système International d'Unités (or SI for short) and is now used in nearly every country in the world. The simplicity of the metric system stems from the fact that there is only one unit of measurement (or base unit) for each type of quantity measured (length, mass, etc.). The three most common base units in the metric system are the metre, gram, and litre. So length, for example, is always measured in metres in the metric system; regardless of whether you are measuring the length of your finger or the distance from Galway to Dublin. To simplify things, very large and very small objects are expressed as multiples of ten of the base unit. For example, rather than saying that distance from Galway to Dublin is 217,000 metres, we can say that it is 217 kilometres. This would be done by adding the prefix "kilo" (meaning 1,000) to the base unit "metre". Metric prefixes can be used with any base unit. For example, a kilometre is 1,000 meters, a kilogram is 1,000 grams, and a kilolitre is 1,000 litres. The metric system is a decimal-based system because it is based on multiples of ten. Any measurement given in one metric unit (e.g., kilogram) can be converted to another metric unit (e.g., gram) simply by moving the decimal place. For example, let's say a friend told you that he weighed 72,500.0 grams. You can convert this to kilograms simply by moving the decimal three places to the left. In other words, your friend weighs 72.5 kilograms. Twenty SI prefixes (below) are available to combine with units of measure. For example, the prefix kilo- denotes a multiple of one thousand, so 1 kilometre equals 1000 metres, 1 kilogram equals 1000 grams, 1 kilowatt equals 1000 watts, and so on. Each SI prefix name has an associated symbol which can be used in combination 6 with the symbols for units of measure. Thus, the "kilo-" symbol, k, can be used to produce km, kg, and kW, (kilometre, kilogram, and kilowatt). Scientific Notation In science, it is common to work with very large and very small numbers. For example, the diameter of a red blood cell is 0.0065 cm, the distance from the earth to the sun is 150,000,000 km, and the number of molecules in 1 g of water is 33,400,000,000,000,000,000,000. It gets cumbersome to work with such long numbers, so measurements such as these are often written using a short hand called scientific notation. Each zero in the numbers above represents a multiple of 10. For example, the number 100 represents 2 multiples of 10 (10 x 10 = 100). In scientific notation, 100 can be written as 1 times 2 multiples of 10: 100 = 1 x 10 x 10 = 1 x 102 (in scientific notation) Scientific notation is a simple way to represent large numbers because the 10's exponent (2 in the previous example) tells you how many places to move the decimal of the coefficient (the one above) to obtain the original number. In our example, the exponent 2 tells us to move the decimal to the right two places to generate the original number: Scientific notation can be used even when the coefficient is a number other than 1. For example: This shorthand can also be used with very small numbers. When scientific notation is used with numbers less than one, the exponent on the 10 is negative, and the decimal is moved to the left, rather than the right. For example: 7 Therefore, using scientific notation, the diameter of a red blood cell is 6.5 x 10-3 cm, the distance from the earth to the sun is 1.5 x 10 8 km and the number of molecules in 1 g of water is 3.34 x 1022. Also note that in scientific notation, the base numeral is always represented as a single digit followed by decimals if necessary. Thus, the number 0.0065 is always represented as 6.5 x 10-3, never as 0.65 x 10-2 or 65 x 10-4. Calculations with Powers of Ten 10 = 101 100 = 10 x 10 = 102 1,000 = 10 x 10 x 10 = 103 1,000,000=10 x 10 x 10 x 10 x 10 x 10 = 106 Multiplying Two Numbers 1,000 x 1,000 = 1,000,000 It would be equivalent to saying that 103 x 103 = 106. Notice that the exponent "6" is just the sum of the other two exponents "3" and "3"! This, in fact is a basic rule whenever you multiply two numbers with the same "base" (in this case, 10). For ANY values a and b, it is always true that 10a x 10b = 10a+b Example (103x105=103+5=108) Now, the exponent can also be a negative number. In this way, we can write 1/10 = 10-1 , 1/100 = 10-2 , and so on. Dividing Two Numbers 100,000/1,000 = 100 — i.e., 105/103= 105-3=102. When dividing two numbers with the same base, the result can be found by subtracting the exponents: 10a/10b = 10a-b. 8 Exercises: 1. Convert each number below into scientific notation (Do not use a calculator). Number Number using Scientific Notation 0.000 934 7 983 000 000 0.000 000 000 820 57 496 x 106 0.000 06 x 101 2. Add, subtract, multiply, or divide the following problems. Express answer in scientific notation (Do not use a calculator). A separate sheet of paper can be used and included in the lab manual, if necessary. Calculation Answer (Use Scientific Notation) (3.21 x 10–3) + (9.21 x 102) (8.1 x 103) + (9.21 x 102) (1.010 1 x 101) – (4.823 x 10–2) (1.209 x 106) x (8.4 x 107) (4.89 x 10–4) ÷ (3.20 x 10–2) Questions 1. How many milligrams (mg) are found in one gram (g)? 2. How many micrograms (µg) are found in one gram (g)? 3. How many nanograms (ng) are found in one gram (g)? 4. How many nanograms (ng) are found in one kilogram (kg)? 9 5. How many nanograms (ng) are found in 10 micrograms (mg)? 6.How many milligrams (mg) are found in 10 kilograms (kg)? Units commonly used in Biochemistry/Biology Prefix Symbol Factor yotta Y 1024 zetta Z 1021 exa E 1018 peta P 1015 tera T 1012 giga G 109 mega M 106 kilo k 103 hecto h 102 deka da 101 deci d 10-1 centi c 10-2 milli m 10-3 micro µ 10-6 nano n 10-9 pico p 10-12 femto f 10-15 atto a 10-18 zepto z 10-21 yocto y 10-24 (The most common unit prefixes range from pico to giga). 10 Lab 1: Making Solutions, pH and How to Use an Autopipette Part A: Making Solutions Learning Outcomes include an understanding of: Percentage Solutions Molar Solutions How to make dilutions Introduction Lab experiments often require preparation of chemical solutions. A number of terms are used and must be understood: Solute - The substance which dissolves in a solution. Solvent - The substance which dissolves another to form a solution. For example, in a sugar and water solution, water is the solvent; sugar is the solute. Solution - A mixture of two or more pure substances. In a solution one pure substance is dissolved in another pure substance homogenously. For example, in a sugar and water solution, the solution has the same concentration throughout, i.e. it is homogenous. Mole - A fundamental unit of mass used by biochemists. A mole is defined as the gram molecular weight of an element or compound, and comprised of exactly 6.023 x 10 23 atoms or molecules (this is called Avogadro's number). 1 mole is 6.02 x 1023 molecules of that substance (Avogadro's number). In biology there are two common ways to represent the concentration of a solute in solution: percent solutions and molar solutions. Percentage Solutions In percentage solutions, the concentration of a solution is often expressed as a "weight/volume percentage". The percentage is calculated from the weight of solute in grams (g), divided by the total volume of solution in millilitres (mL): (Mass(g) / Volume(mL))x100 = % A 1% solution would therefore have 1 g of solute dissolved in a final volume of 100 mL of solution. This would be labelled as a weight/volume (w/v) percentage solution. 11 When using liquid reagents the percent concentration is based upon volume per volume, and is similarly calculated as % concentration x volume needed = volume of reagent to use. Moles and Molar Solutions (unit = M = moles/L) Sometimes it may be more efficient to use molarity when calculating concentrations. A mole is defined as the gram molecular weight of an element or compound. The mass (g) of one mole of an element is called its molecular weight (MW). When working with compounds, the mass of one mole of the compound is called the formula weight (FW). The distinction between MW and FW is not always simple, however, and the terms are routinely used interchangeably in practice. Formula (or molecular) weight is always given as part of the information on the label of a chemical bottle. The number of moles in an arbitrary mass of a dry reagent can be calculated as: No. of moles = weight (g)/ molecular weight (g) Molarity is the unit used to describe the number of moles of a chemical or compounds in one litre (L) of solution and is thus a unit of concentration. By this definition, a 1.0 Molar (1.0 M) solution is equivalent to one formula weight (FW = g/mole) of a compound dissolved in 1 litre (1.0 L) of solvent (often water). E.g. Sodium Chloride (NaCl) has a FW of 58.44. Therefore, a 1M solution of NaCl in water contains 58.44 grams of NaCl in 1 litre of water. Example 1: To prepare a litre of a simple molar solution from a dry reagent Multiply the formula weight (or MW) by the desired molarity to determine how many grams of reagent to use: Chemical FW = 194.3 g/mole; to make 0.15 M solution use 194.3 g/mole * 0.15 moles/L = 29.145 g/L Example 2: To prepare a specific volume of a specific molar solution from a dry reagent A chemical has a FW of 180 g/mole and you need 25 ml (0.025 L) of 0.15 M (M = moles/L) solution. How many grams of the chemical must be dissolved in 25 ml water to make this solution? No. of grams/desired volume (L) = desired molarity (mole/L) * FW (g/mole) by algebraic rearrangement, No. of grams = desired volume (L) * desired molarity (mole/L) * FW (g/mole) No of grams = 0.025 L * 0.15 mole/L * 180 g/mole After cancelling the units, 12 No. of grams = 0.675 g So, you need 0.675 g/25 ml Simple Dilution (Dilution Factor Method Based on Ratios) A simple dilution is one in which a unit volume of a liquid material of interest is combined with an appropriate volume of a solvent liquid to achieve the desired concentration. The dilution factor is the total number of unit volumes in which your material will be dissolved. The diluted material must then be thoroughly mixed to achieve the true dilution. For example, a 1:5 dilution (verbalize as "1 to 5" dilution) entails combining 1 unit volume of solute (the material to be diluted) + 4 unit volumes of the solvent medium (hence, 1 + 4 = 5 = dilution factor). Example 1: Frozen orange juice concentrate is usually diluted with 4 additional cans of cold water (the dilution solvent) giving a dilution factor of 5, i.e., the orange concentrate represents one unit volume to which you have added 4 more cans (same unit volumes) of water. So the orange concentrate is now distributed through 5 unit volumes. This would be called a 1:5 dilution, and the OJ is now 1/5 as concentrated as it was originally. So, in a simple dilution, add one less unit volume of solvent than the desired dilution factor value. Example 2: Suppose you must prepare 400 ml of a disinfectant that requires 1:8 dilution from a concentrated stock solution with water. Divide the volume needed by the dilution factor (400 ml / 8 = 50 ml) to determine the unit volume. The dilution is then made with 50 ml concentrated disinfectant + 350 ml water. Making Solutions Experiment 1 (a) 0.1% w/v Sodium Chloride (NaCl) Solution A 1% NaCl solution has 1 gram of sodium chloride dissolved in 100 ml of solution. Therefore, a 0.1% NaCl solution has 0.1 grams of sodium chloride dissolved in 100 ml of solution. Procedure: Weigh 0.1 g of sodium chloride. Put it into a 100 ml beaker containing about 80ml of water. Once the sodium chloride has dissolved completely (swirl if necessary), pour the solution into a 100ml volumetric flask, add water to bring the volume up to the 100 ml mark (bottom of meniscus touching the line). Caution: You cannot simply measure 100ml of water (H2O) and add 0.1g of sodium chloride. This will introduce error because adding the solid will change the final volume of the solution and throw off the final percentage. 13 Experiment 1 (b): 1% v/v Ethanol Solution When the solute is a liquid, it is sometimes convenient to express the solution concentration as a volume percent. A 1% ethanol solution has 1 ml of ethanol in 100 ml of solution. Procedure: Pipette exactly 1 ml of ethanol into a 100 ml volumetric flask, then, add water (H2O) to fill the volumetric flask to the 100 ml mark indicated on the flask. Experiment 1(c): 0.05 M (50mM) Sucrose in H2O solution The formula for molarity (M) is: moles of solute / 1 litre of solution or gram-molecular masses of solute / 1 litre of solution. The FW/MW of sucrose is 342. A 1M sucrose solution contains 342 g of sucrose in 1000 ml (1 L). Therefore, a 0.05 M sucrose solution requires 17.1g of sucrose in 1000ml. Procedure: Weigh 17.1 g of sucrose. Put into a 200ml beaker and add water to dissolve (swirl if necessary). Pour the solution into a 1000ml volumetric flask and add water to bring the volume to 1000 ml. Part B: pH and How to Measure pH Learning Outcomes  An understanding of pH  How to calibrate a pH meter  How to measure the pH of solutions Introduction There are millions of chemical substances in the world. Some of them have acidic properties, others, basic properties. Acids are substances which free hydrogen ions (H+), when they are mixed with water. Bases are substances which free hydroxide ions (OH-) when they are mixed with water. (This freeing of ions is called dissociation in both cases). Free hydroxide ions react with the hydrogen ions producing water molecules: H+ + OH- = H2O. In this way, bases diminish the concentration of hydrogen ions. A solution rich in hydrogen ions is acidic, a solution poor in hydrogen ions is basic. Some acids dissociate only in part and they are called weak acids; others dissociate completely, freeing large amounts of hydrogen ions, and they are called strong acids. In the same way, the bases can be stronger or weaker. Diluted acids and bases are less concentrated and less aggressive in their actions. The acidic or basic degree of substances is measured in pH units. The scale used spans 14 from 0 to 14. Substances with pH lower than 7 are considered acids, those with pH equal to 7 are considered neutral, and those with pH higher than 7 are considered bases. Substances with low pH are very acidic, while those with high pH are highly basic. Concentrated acidic and basic substances are very corrosive and dangerous. pH is the measure of the concentration of hydrogen ions in a solution. As this concentration can extend over several orders of magnitude, it is convenient to express it by means of logarithms of base ten. As this concentration is always less than one, its logarithm always has the minus sign. To avoid having to always write the minus sign, it has been agreed to write this value with the positive sign. (This is the same as using the logarithm of the reciprocal of the hydrogen ion concentration). So, the pH is the logarithm of the concentration of hydrogen ions, with the sign changed: pH = - log[H+]. Thus, when pH has low values, the concentration of hydrogen ions is high. Distilled water has pH = 7. So how it is possible that distilled water has free hydrogen ions? Their presence is due to the casual dissociation of some water molecules because of the thermal agitation (H2O ↔H+ + OH-). Immediately after, these ions recombine themselves, but other molecules dissociate themselves, thus keeping a constant equilibrium of a certain concentration of dissociated molecules. Measuring pH There are substances which have the property of changing their colour when they come in contact with an acidic or basic environment. These substances are called pH indicators. Usually, they are used in solution e.g. phenolphthalein and bromothymol blue. Special papers which have been soaked with indicators may be used to get a rough estimate of pH. These papers change colour when they are immersed in acidic or basic liquids. This is the case of the well-known litmus paper. Generally in laboratories, pH is measured more accurately using a pH meter. A pH meter measures the pH of a solution utilizing a glass electrode. The electrode is made of very thin glass that allows H+ ions to pass through it. The meter measures electrical potential and converts this data into a pH reading for a sample. Demonstration: Calibration of a pH Meter All pH meters need to be calibrated on a regular basis; as time passes by and with frequent use a pH meter it will lose calibration. The method of calibration will depend on the type of pH meter used. pH meters require calibration with solutions of two different pHs e.g. pH 7 (neutral) and pH 4 (acidic). If a pH meter is used every day it is advisable to calibrate it once a week. To calibrate a pH meter, you need three clean glass/plastic containers that can hold sufficient solution to immerse the pH probe. One container will hold water for rinsing the probe, another container for the pH 7 solution and the last container for the pH 4 solution. Calibration of a Hanna pH Meter 1. Rinse probe in water and briefly shake off excess. 2. Switch the meter on. 3. Place probe into pH 7 solution and allow the pH reading to settle. 4. Adjust pH 7 screw on the top side of the meter until the meter reads 7. 5. Rinse probe again in water and shake. 15  6. Place probe into pH 4 solution and allow the pH reading to settle. 7. Adjust the pH 4/10 screw on the top side of the meter until the meter reads 4. The pH meter is now calibrated. Reagents and equipment required to measure the pH of some common solutions (i) pH Meter (ii) Various solutions including e.g. distilled water, tap water, lemon juice, cola, wine etc. (iii) Tubes with caps Experiment 2: pH of some Common Solutions 1. Measure the pH of each of the 5mL aliquots of common solutions provided. 2. Remember to rinse and dry the pH probe between measurements. 3. Record the pH of each solution in the table provided below and describe each solution in terms of being acidic, alkaline or neutral pH. Solution pH Description Distilled Water Tap Water Vinegar Lemon Juice Cola Questions: 1. Explain what is meant by a 10% w/v sodium chloride solution. 2. What is meant by the term mole? 3. Explain how to make a 2M sodium chloride solution. 4. What is meant by a 1 in 5 dilution? Part C: How to Deliver an Accurate Volume Using an Autopipette 16 Learning Outcomes: How to use an autopipette How to calibrate an autopipette Introduction The equipment you select to measure and deliver a volume of liquid depends on (a) the volume, (b) the accuracy, and (c) the number of times the measurement/delivery is required. In addition, certain types of liquids may present problems, which may influence the choice of equipment, e.g. high-viscosity liquids, organic solvents, solutions that froth, or suspensions that sediment. Some of the options are given in the table below. Methods to dispense liquids and criteria for selection of a method Method Optimum Accuracy Usefulness for repeated volume range measurements Graduated 5-2000 mL Medium Reasonably good cylinder 5-2000 mL High Reasonably good Volumetric flask 25-5000 mL Very low Reasonable good Conical flask/Beaker 30 L –2 mL Low to Very good Pasteur pipette Medium 1 L –10 mL Very good Automatic Very high* pipette/glass Any volume Reasonably good pipette Very high* Weighing *Automatic pipettes must be calibrated regularly and accurately (see below). Important Points When Using an Autopipette 1. Select the pipette that is suitable for the volume range you are using, as most are accurate only over specified ranges. Never attempt to set the volume above or below the maximum and minimum limits, respectively - otherwise the pipette may be damaged! The maximum volume is generally given on the button on the end of the metal plunger, e.g. P10 (10 L or 0.01 mL), P20 (20 L or 0.02 mL) or P1000 (1000 L or 1.0 mL) 2. Using a twisting motion, set the pipette to deliver the desired volume. Each pipette will normally have a volume scale dial. If you dial 100 on a P1000 pipette, the volume is set at 1000 L. If you dial 100 on a P20 autopipette, the volume is set at 10.0 L (the red digit is 0.1 L). 3. Fit a clean disposable tip on the end of the pipette - make sure you are using the correct type (blue for a P1000, and yellow for a P200 or P20). Take care to fit the tip correctly, as liquid will leak from the tip otherwise. NEVER use a 17  pipette without its disposable tip! NEVER leave a pipette on its side with the tip full of liquid! 4. Holding the pipette vertically, press the metal plunger down slowly until you meet the first point of resistance. Place the tip in the liquid and gradually and evenly release the metal plunger. Make sure that the tip of the disposable pipette tip is in the liquid to avoid air bubbles. Wait a moment or two to ensure that all of the liquid has been drawn up, then, withdraw the pipette tip from the liquid. 5. If you draw the liquid up too quickly or in a jerking manner, you run the risk of drawing liquid up into the barrel of the pipette. If this happens, seek assistance from your demonstrator or class supervisor/technical officers, as the barrel must be cleaned before further use. 6. Deliver the volume of liquid in the tip, by placing the end of the tip at a slight angle against the wall of a vessel (e.g. test-tube or flask). Press the plunger slowly to release the liquid. Allow any residual liquid to drain from the walls of the tip and press the plunger to release this residual volume. 7. Remove the pipette and eject the tip into the yellow sharps container on your bench and not randomly anywhere on your bench! At the end of the practical, empty the contents of the yellow sharps container into the glass/sharps bin at the top of the laboratory. 8. N.B. – important: When pressing the metal plunger down to take an aliquot of liquid, you should only press as far as the first point of resistance. If you go beyond this point, you will take up an additional volume of liquid. If this happens, you should only press the plunger to the first end-point when releasing the liquid, and not attempt to press beyond this point. Otherwise you will dispense an excess of your solution. 9. Before starting any experiment it is essential to check the accuracy and precision of your autopipettes. 18 Experiment 3: Calibration of an P1000 Autopipette 1. Turn on the balance. Place a small empty vessel (e.g. weigh boat or small plastic beaker) on the balance pan, and press the tare bar to zero. 2. Using the technique outlined above, calibrate your P1000 first, by setting the volume scale dial to deliver 1.0 mL (1000 L). Withdraw 1.0 mL of liquid (H2O), and deliver to the vessel on the balance pan – take care not to spill any liquid on the balance pan. Record the weight. Press the tare bar to zero. Repeat the procedure with a fresh aliquot of liquid. Note the weight and zero again. Repeat for a further 5 liquid deliveries. 3. Calculate the mean, standard deviation and the coefficient of variation for the set of data. Comment on the accuracy and precision (See how to calculate below). 19 Accuracy and precision Accuracy and precision are essential in conducting experimental and operational procedures. Quality of performance depends on both parameters; therefore a clear understanding of both terms is required. Illustration of the concepts of accuracy and precision by analogy with results from a target shooting contest. Note that the imprecise but accurate result is accurate only when a mean result is estimated from a large number of individual results, and that the result, which is precise but inaccurate, may also be termed biased. Good accuracy for the analyst means obtaining the correct result, and is expressed as a percentage, with 100% being desirable. Accuracy (%) = actual result x 100 correct result Good precision for the analyst means being able to obtain the same result on repeated analysis, and is expressed as a percentage, with 0% being desirable. The statistic used is the percent standard deviation or coefficient of variation (CV), which facilitates comparison of measurements of precision. CV (%) = standard deviation x 100 mean The standard deviation (sd) is the statistic used to express the degree to which a repeated result varies. It is related to the mean deviation (i.e. the average of all the differences between the mean and the individual results) but is much more useful. sd = [x1 – mean]2 + [x2 – mean]2 +… [xn – mean]2 n-1 x = result or one value in your set of data mean = average off all values x in your set of data n = the number of values x in your set of data 20 Lab 2: Proteins and Nucleic Acids Part A: Proteins Learning Outcomes include an understanding of: Protein Structure and Function Methods for Assaying Proteins Introduction Most proteins are linear polymers built from series of up to 20 different L-α-amino acids. All amino acids possess common structural features, including an α-carbon to which an amino group, a carboxyl group, a Hydrogen atom (H) as well as a variable side chain. The amino acids in a polypeptide chain are linked by peptide bonds. Once linked in the protein chain, an individual amino acid is called a residue. The end of the protein with a free carboxyl group is known as the carboxy terminus (C-terminus) , whereas the end with a free amino group is known as the amino terminus (N-terminus). Proteins are assembled from amino acids using information encoded in genes. Each protein has its own unique amino acid sequence that is specified by the nucleotide sequence of the gene encoding this protein. The genetic code is a set of three- nucleotide sets called codons and each three-nucleotide combination designates an amino acid, for example AUG (adenine-uracil-guanine) is the code for methionine. Because DNA contains four nucleotides, the total number of possible codons is 64; hence, there is some redundancy in the genetic code, with some amino acids specified by more than one codon. Biochemists often refer to four distinct levels of protein structure (diagram below):  Primary structure: the amino acid sequence i.e. amino acids linked by peptide bonds.  Secondary structure: regularly repeating local structures stabilized by hydrogen bonds. The most common examples are α helices and β pleated sheets. Because secondary structures are local, many regions of different secondary structure can be present in the same protein molecule.  Tertiary structure: the overall shape of a single protein molecule; the spatial relationship of the secondary structures to one another. Tertiary structure is generally stabilized by non-local interactions, most commonly the formation of a hydrophobic core, but also through ionic bonds, hydrogen bonds, disulfide bonds, and even post-translational modifications. The term "tertiary structure" is often used as synonymous with the term fold. The tertiary structure is what controls the basic function of the protein.  Quaternary structure: the structure formed by several protein molecules (polypeptide chains), usually called protein subunits in this context, which function as a single protein complex. 21 Proteins can be informally divided into two main classes, which correlate with typical tertiary structures: globular proteins and fibrous proteins. Almost all globular proteins are soluble and many are enzymes. Fibrous proteins are often structural, such as collagen, the major component of connective tissue, or keratin, the protein component of hair and nails. Good food sources of protein include meat, fish, eggs, dairy products. Protein Assays Protein assays determine the amount of protein in an unknown solution. Proteins are polymers of amino acids, where amino acid units are joined by peptide bonds which have partial double bond character and absorb light. Proteins also may contain some amino acids with aromatic rings in the side chains. Proteins may be assayed using light absorbance or spectrophotometry of light with a wavelength of 280 nm. Techniques for determining protein concentration depend on using a standard protein to calibrate measurements for accuracy. All protein assay techniques are prone to inaccuracy due to a variety of variables that may affect absorbency (i.e. interfering substances). There are several types of protein assay used in Biochemistry: The Biuret Assay uses a Cu2+ peptide bond complex. The Biuret method is not highly sensitive but is the most linear because the colour produced depends on a direct complex between peptide bonds of the protein and Cu2+ ion. The Biuret assay is the method used in the experiment below. The Lowry Method relies on the presence of aromatic amino acids in the protein. It involves the formation of a heavy metal complex with the aromatic amino acids. First a peptide bond complex is formed which is then enhanced by a phospho-molybodate complex with the aromatic amino acids. 22 The Bradford Assay involves a dye reaction with amino group side chains. It is based on an absorbance shift in a dye known as Coomassie Blue when bound to arginine and aromatic amino acid residues found in protein. Unlike other protein assays, the Bradford Protein Assay is less sensitive to interference by various chemicals that could exist in protein samples Reagents and Equipment for Detecting Proteins: Biuret Reagent (N.B. Wear gloves when handling Biuret. Skin and nails contain protein!) Bovine Serum Albumin (BSA) Protein Standard Milk Albumin (Egg White) Distilled H2O Test tubes with lids Autopipette (1000 µL) Blue tips Experiment 2(a) Testing for Proteins - Biuret Assay 1. Pipette 2 ml of BSA protein standard solution to a test tube. 2. Pipette 2 ml of water to another test tube. 3. Pipette 2 ml of Patient A mock urine sample to a third test tube. 4. Pipette 2 ml of Patient B mock urine sample to a fourth test tube. 5. Pipette 2 ml of Biuret reagent to each of the test tubes above. 6. Incubate tubes at room temperature for 5 minutes. 7. Write down your observations and discuss in the box provided below. 23 Questions: What is meant by protein primary structure? What is meant by protein secondary structure? What is meant by protein tertiary structure? Part B: Nucleic Acids Learning Outcomes include an understanding of:  Nucleic Acid Structure and Function  How to purify DNA by alcohol precipitation Introduction Nucleic acids are linear, unbranched polymers of nucleotides. Each nucleotide consists of three components: a nitrogenous base, which is either a purine (adenine and guanine) or a pyrimidine (thymine and cytosine); a pentose sugar; and a phosphate group. There are two types of nucleic acid: Ribonucleic acid (RNA) and deoxyribonucleic acid (DNA). 24 DNA and RNA differ in the structure of the sugar in their nucleotides - DNA contains 2-deoxyribose while RNA contains ribose (the only difference is the presence of a hydroxyl group). Also, the nitrogenous bases found in the two nucleic acid types are different: adenine, cytosine, and guanine are found in both RNA and DNA, while thymine only occurs in DNA and uracil only occurs in RNA. The sugars and phosphates in nucleic acids are connected to each other in an alternating chain, linked by phosphodiester bonds. Ribonucleic acid, or RNA, is a nucleic acid polymer consisting of nucleotide monomers, which plays several important roles in the processes of transcribing genetic information from deoxyribonucleic acid (DNA) into proteins. RNA acts as a messenger between DNA and the protein synthesis complexes known as ribosomes, forms vital portions of ribosomes, and serves as an essential carrier molecule for amino acids to be used in protein synthesis. The three types of RNA include tRNA (transfer), mRNA (messenger) and rRNA (ribosomal). Deoxyribonucleic acid is a nucleic acid that contains the genetic instructions used in the development and functioning of all known living organisms. The main role of DNA molecules is the long-term storage of information and DNA is often compared to a set of blueprints, since it contains the instructions needed to construct other components of cells, such as proteins and RNA molecules. The DNA segments that carry this genetic information are called genes, but other DNA sequences have structural purposes, or are involved in regulating the use of this genetic information. DNA generally exists as a double stranded molecule held together by hydrogen bonding between complementary bases. Adenine bonds to thymine and cytosine bonds to guanine. The bonds between the bases are Hydrogen bonds. The DNA chains coil around each other, forming the DNA double helix (diagram overleaf). Extracting DNA from cells is not an easy task; the cell walls and nuclear membranes must be lysed (broken) so that the DNA molecules can be dissolved into a solvent. The DNA must then be separated from cell debris, substructures and other molecules. Cells also contain a variety of enzymes, some of which attack and destroy nucleic acids. The DNA must be protected from these nucleases as it is being isolated. Human DNA is generally prepared from peripheral blood samples or from tissue samples e.g. skin. This practical demonstrates a final stage in DNA isolation. The addition of alcohol to a solution will precipitate DNA, allowing it to be spooled with a glass rod. Ethanol precipitation is a commonly used technique for concentrating and de-salting nucleic acid (DNA or RNA) preparations in aqueous solution. The basic procedure is that salt and ethanol are added to the aqueous solution, which forces the nucleic acid to precipitate out of solution. 25 Water is a polar molecule – it has a partial negative charge near the oxygen atom due the unshared pairs of electrons, and partial positive charges near the hydrogen atoms. Because of these charges, polar molecules, like DNA or RNA, can interact electrostatically with the water molecules, allowing them to easily dissolve in water. Polar molecules can therefore be described as hydrophilic and non-polar molecules, which can’t easily interact with water molecules, are hydrophobic. Nucleic acids are hydrophilic due to the negatively charged phosphate (PO3-) groups along the sugar phosphate backbone. The role of the salt in the protocol is to neutralize the charges on the sugar phosphate backbone. A commonly used salt is sodium acetate. In solution, sodium acetate breaks up into Na+ and [CH3COO]-. The positively charged sodium ions neutralize the negative charge on the PO3- groups on the nucleic acids, making the molecule far less hydrophilic, and therefore much less soluble in water. The electrostatic attraction between the Na+ ions in solution and the PO3- ions are dictated by Coulomb’s Law, which is affected by the dielectric constant of the solution. Water has a high dielectric constant, which makes it fairly difficult for the Na+ and PO3- to come together. Ethanol on the other hand has a much lower dielectric constant, making it much easier for Na+ to interact with the PO3-, shield its charge and make the nucleic acid less hydrophilic, causing it to drop out of solution. Reagents and Equipment required for DNA Precipitation (i) Deoxyribonucleic Acid (Salmon Sperm DNA) (ii) Sodium Acetate (iii) Cold Ethanol (iv) Plastic Test Tube with lid (DNA sticks to glass) (v) Glass Rod Experiment 2(b): Ethanol Precipitation and Spooling of Salmon Sperm DNA 1. Pipette 0.5 ml of DNA solution into a plastic test tube. 2. Slowly pipette 0.5 ml 3M sodium acetate solution. 3. Mix gently by swirling. 4. Slowly, pipette approximately 2 ml of absolute ethanol. 5. Mix by inverting gently after tightly replacing top. 6. Observe DNA precipitation. 7. Gently spool DNA with glass rod provided. 8. Write down your observations in the box overleaf. 26 Questions: 1. What is meant by the term nucleotide? 2. What bases in DNA are purines? 3. What bases in DNA are pyrimidines. 4. What are the main differences in structure between DNA and RNA? 27 Part C: Agarose Gel Electrophoresis Learning Outcomes include an understanding of: How electrophoresis works. How to prepare an agarose gel. How to load an agarose gel. Introduction Agarose gel electrophoresis is an easy way to separate DNA fragments by their sizes and visualize them. It is a common diagnostic procedure used in molecular biological labs. The technique of electrophoresis is based on the fact that DNA is negatively charged at neutral pH due to its phosphate backbone. For this reason, when an electrical potential is placed on the DNA it will move toward the positive electrode of a gel apparatus. The rate at which the DNA will move toward the positive electrode is slowed by making the DNA move through an agarose gel. This is a buffer solution (which maintains the proper pH and salt concentration) with 0.75% to 2.0% agarose added. The agarose forms a porous lattice in the buffer solution and the DNA must slip through the holes in the lattice in order to move toward the positive pole. This slows the molecule down. Larger molecules will be slowed down more than smaller molecules, since the smaller molecules can fit through the holes As a result, a mixture of large and small fragments of DNA that has been run through an agarose gel will be separated by size. Other charged molecules such as proteins and dyes can be subjected to electrophoresis.. 28 Reagents and Equipment for Gel Electrophoresis of Dyes (i) 2% Agarose Gel (ii) Tris – Acetate – EDTA Buffer (TAE) (iii) Gel Apparatus (iv) P20 Autopipette (v) Various red, green, blue, yellow food dyes in 10% sucrose solution (iv) Mix of red, green, blue and yellow dyes in 10% Sucrose solution Experiment 2(c) – Agarose Gel Electrophoresis of Dyes (1) Set a P20 pipette to 10 µL. (2) Pipette 10 µL of one dye sample into one well of the agarose gel (the demonstrators will help you with this). Make a note of which lane you have loaded. (3) After the gel is loaded, a demonstrator will connect the apparatus to a power pack and a voltage of 100 volts will be applied to the gel for 20 minutes. (4) Draw a diagram of the gel and explain why the dyes have moved various distances (using the box overleaf). 29 Lab 3 – Lipids and Carbohydrates Part A: Lipids Learning Outcomes include an understanding of: The nature of lipids How to test for lipids Introduction Lipids as group of molecules includes fats and oils, waxes, phospholipids, steroids (like cholesterol), and some other related compounds. All lipids are hydrophobic. Fats and oils are made from two kinds of molecules: glycerol (a type of alcohol with a hydroxyl group on each of its three carbons) (diagram to left) and three fatty acids. Since there are three fatty acids attached, these are known as triglycerides (diagram below). The main distinction between fats and oils is whether they are solid or liquid at room temperature. This is based on differences in the structures of the fatty acids they contain. The “tail” of a fatty acid is a long hydrocarbon chain, making it hydrophobic. The “head” of the molecule is a carboxyl group which is hydrophilic. The terms saturated, mono-unsaturated, and poly-unsaturated refer to the number of hydrogens attached to the hydrocarbon tails of the fatty acids as compared to the number of double bonds between carbon atoms in the tail. Fats, which are mostly from animal sources, have single bonds between the carbons in their fatty acid tails, thus all the carbons are also bonded to the maximum number of hydrogens possible. Since the fatty acids in these triglycerides contain the maximum possible amount of hydrogens, are called saturated fats. The hydrocarbon chains in these fatty acids are fairly straight and can pack closely together, making these fats solid at room temperature. Oils, mostly from plant sources, have some double bonds between some of the carbons in the hydrocarbon tail, causing bends or “kinks” in the shape of the molecules. These oils are called 30 unsaturated fats because the double bond(s) mean that they are not saturated with hydrogen. The unsaturated fats are unable to pack as closely together because of kinks made by the double bonds, and this means that they are liquid at room temperature. Unsaturated fats in the diet are considered to be “healthier” than the saturated fats. Lipids function as long-term energy storage. One gram of fat stores more than twice as much energy as one gram of carbohydrates. Lipids are stored in adipose tissue in humans. Phospholipids Phospholipids are made from glycerol, two fatty acids, and (in place of the third fatty acid) a phosphate group with some other molecule attached to its other end. The hydrocarbon tails of the fatty acids are still hydrophobic, but the phosphate group end of the molecule is hydrophilic. This means that phospholipids are amphipathic i.e. soluble in both water and oil. Cell membranes are made mostly of phospholipids arranged in a double layer with the tails from both layers “inside” (facing toward each other) and the heads facing “out” (toward the watery environment) on both surfaces. Steroids The general structure of cholesterol consists of three six-membered rings side-by- side and one five-membered ring (diagram below). The central core of this molecule, consisting of four fused rings, is shared by all steroids, including hormones such as oestrogen, testosterone and progesterone. In the various types of steroids, various other groups/molecules are attached around the edges. Food Sources of Lipids Lipids are found in meat, dairy products such as milk and butter, eggs, seeds and nuts etc. Testing for Lipids The two basic tests for lipids which are used in this class are the Emulsion test and the Grease Spot test. In the Emulsion test, the lipid sample to be suspended in ethanol, allowing lipids present to dissolve (lipids are soluble in alcohols). The 31 liquid (alcohol with dissolved fat) is then decanted into water. Since lipids do not dissolve in water while ethanol does, when the ethanol is diluted, it falls out of the solution to give a cloudy white emulsion. The Grease Spot test can also be used identify the presence of lipids in a sample as brown paper becomes translucent in the presence of lipids. Another test is the Sudan Red test. Sudan Red is a lipid soluble dye and stains lipids red. Sudan red may show the amount and the location of lipids. Reagents and Equipment for Detecting Lipids i. Aliquots of various oils e.g. olive oil and vegetable oil, skim milk and full fat milk, egg white and egg yolk solutions ii. Ethanol iii. Distilled H2O iv. Squares of brown paper and cotton swabs v. Test tubes with lids Experiment 3(a) Testing for the presence of lipids: Emulsion Test for Lipids 1. Pipette 1ml of any oil and 1ml of ethanol into a capped tube (tube 1). 2. Replace the lipid firmly and shake the tube vigorously for one minute. 3. Pipette 2ml of distilled H20 into another tube (tube 2). 4. Carefully, pour the contents of tube 1 into tube 2. Replace the lid. 5. Write down your observations. 6. Comment on your results. 32  Experiment 3(b) Testing for the presence of lipids: Grease spot test 1. Draw four squares onto the brown paper provided. 2. Then, use a cotton swab to put samples of three lipids of your choice onto the brown paper. 3. Use a cotton swab to put a water sample onto the brown paper, as a control. 4. Wipe off excess oil/fat and let sit for few minutes to dry. 5. Once dry, the fats will have left a translucent spot behind. This can best be seen when you hold the paper up to a light source. 6. Write down your observations from the experiment and comment on the results in the box below Questions: 1. What is a saturated fat? 2. Explain what is meant by the term triglyceride. 3. What molecule is the precursor of steroids such as steroid hormones? 33 Part B: Carbohydrates Learning Outcomes include an understanding of: The nature of carbohydrates (monosaccharides, disaccharides and polysaccharides) How to test for monosaccharides (reducing sugars) How to test for polysaccharides Introduction There are three types of carbohydrate: monosaccharides (simple sugars); dissacharides and polysaccharides. Monosaccharides or simple sugars such as glucose and fructose function as energy sources in cells during cellular respiration. They are also used to build cell structures and other organic molecules within the cells. Glucose is the primary form of sugar stored in the human body for energy. Fructose is the main sugar found in most fruits. Both glucose and fructose have the same chemical formula (C6H12O6); however, they have different structures, as shown (note: the carbon atoms that sit in the "corners" of the rings are not labelled. Glucose Fructose Disaccharides have two sugar units bonded together by a glycosidic bond. For example, common table sugar is sucrose, a disaccharide that consists of a glucose unit bonded to a fructose unit. Sucrose 34 Polysaccharides or complex carbohydrates are long chains of monosaccharides bonded together by glycosidic bonds. Plants store excess glucose in the form of starch, a polysaccharide composed of long chains of glucose. Polysaccharides or starches are found in potatoes, rice, wheat, corn, bananas, peas, beans, lentils, and other tubers, seeds and fruits of plants. Animals (and humans) store excess glucose in the form of glycogen in the liver and muscles. Between meals the liver breaks down glycogen to glucose and releases it into the blood stream to supply glucose to cells as required. Other important polysaccharides are cellulose and chitin. Cellulose makes up the cell wall of plants. Chitin provides structure to fungi and the exoskeleton of arthropods. The potato contains the complex carbohydrate starch. Starch is a polymer of the monosaccharide glucose. Starch (n is the number of repeating glucose units and ranges in the 1,000's) Starch is the principal polysaccharide used by plants to store glucose for later use as energy. When humans eat starch, an enzyme called amylase that occurs in saliva and in the intestines breaks the bonds between the repeating glucose units, thus allowing the sugar to be absorbed into the bloodstream. Once absorbed into the bloodstream, the human body distributes glucose to the areas where it is needed for energy or stores it as its own special polymer – glycogen. Glycogen, another polymer of glucose, is the polysaccharide used by animals to store energy. Excess glucose is bonded together to form glycogen molecules, which the animal stores in the liver and muscle tissue as an "instant" source of energy. Both starch and glycogen are polymers of glucose; however, starch is a long, straight chain of glucose units, whereas glycogen is a branched chain of glucose units. Another important polysaccharide is cellulose. Cellulose is yet another polymer of the monosaccharide glucose. Cellulose, also known as plant fibre, cannot be digested by humans, therefore cellulose passes through the digestive tract without being absorbed into the body. Some animals, such as cows and termites, contain bacteria in their digestive tract that help them to digest cellulose. Cellulose is a 35 relatively stiff material, and in plants it is used as a structural molecule to add support to the leaves, stem, and other plant parts. Reducing Sugars A reducing sugar is any sugar that, in a solution, has an aldehyde or a keto group. This allows the sugar to act as a reducing agent, for example, in Benedict's reaction. Reducing monosaccharides include glucose, fructose, glyceraldehyde and galactose. Many disaccharides, like lactose and maltose also have a reducing form, as one of the two units may have an open-chain with an aldehyde group. However, sucrose, in which the anomeric carbons of the two units are linked together, is a non- reducing disaccharide. Food Sources of Carbohydrates Examples of single sugars from foods include:  Fructose (found in fruits)  Galactose (found in milk products) Disaccharides include:  Lactose (found in milk)  Maltose (found in certain vegetables and in beer)  Sucrose (table sugar) Polysaccharides/complex carbohydrates, often referred to as "starchy" foods, include: legumes, starchy vegetables, cereals and breads. Testing for Carbohydrates Lugol's reagent (iodine solution) changes from yellowish-brown to dark purple/black in the presence of polysaccharide. Benedict's solution (also called Benedict's reagent or Benedict's test) is a chemical reagent named after an American chemist, Stanley Rossiter Benedict. Benedict’s solution is used to test for simple carbohydrates (monosaccharides and some disaccharides). Benedict's solution is a blue coloured liquid that contains copper ions. When Benedict's solution and simple carbohydrates are heated, the solution changes to orange red/ brick red. This reaction is caused by the reducing property of simple carbohydrates. The copper (II) ions (Cu2+) in the Benedict's solution are reduced to copper (I) ions (Cu+), which causes the colour change. Sometimes a brick red solid, copper oxide, precipitates out of the solution and collects at the bottom of the test tube. The intensity of the colour depends on the concentration of glucose present in the sample. Complex carbohydrates such as starches do not give a positive with the Benedict's test unless they are broken down through heating or digestion. Sucrose or table sugar (disaccharide) is a non-reducing sugar and does also not react with the iodine or with the Benedict Reagent. Sugar needs to be broken down into its components, glucose and fructose, to give a positive with Benedict’ solution but the starch test will still be negative. 36 Reagents and equipment for detecting sugars (i) Glucose solution, (ii) Sucrose ( table sugar) solution (iii) Starch (potato, pasta or bread) solution (iv) Lugol’s (iodine)solution (v) Benedict’s solution (vi) Water bath (vii) Test tubes with lids Experiment 3(c) Test for monosaccharides: Benedict’sTest 1. Pipette 2 ml of glucose solution into a test tube. 2. Pipette 2 ml of (sucrose) table sugar solution into a second test tube 3. Pipette 2 ml of starch (pasta, potato or bread) solution into a third test tube. 4. Pipette 1 ml of the Benedict’s solution to each. 5. Incubate the tubes in a 50 °C water bath for 10 minutes. 6. Remove test tubes from water bath (care must be taken – the tubes may be hot). 7. Write down your observations concerning the test samples and discuss below. Experiment 3(d) Test for Starch (Polysaccharide): Iodine Solution 1. Pipette 2 ml of glucose solution into a test tube 2. Pipette 2 ml of (sucrose) table sugar solution into a test tube 3. Pipette 2 ml of starch solution into a test tube. 4. Pipette 2 ml of potato or pasta or bread solution into a test tube. 5. Add 3-4 drops of iodine to each tube. A bluish black colour indicates a positive test for starch. 6. Write down observations for each test and discuss below. 37 Questions: 1. What is a monosaccharide? Give three examples. 2. What is disaccharide? Give three examples. 3. What is a polysaccharide? Give three examples. 4. What is a polymer? 5. What is a reducing sugar? 6. How does Benedict’s solution allow detection of reducing sugars? 38 Notes