Jul 24, 2013

ANALYSIS OF VEGETABLES AND FRUIT JUICES

AIM

To analyse some fruits & vegetables juice for the contents present in them.

INTRODUCTION

Fruits and vegetable are always a part of balanced diet. That means fruits vegetables provide our body the essential nutrients, i.e. Carbohydrates, proteins, vitamins and minerals. Again their presence in these is being indicated by some of our general observations, like -freshly cut apples become reddish black after some time. Explanation for it is that iron present in apple gets oxidixed to iron oxide. So, we can conclude that fruits and vegetables contain complex organic compounds, for e.g., anthocin, chlorophyll, esters(flavouring compounds), carbohydrates, vitamins and can be tested in any fruits or vegetable by extracting out its juice and then subtracting it to various tests which are for detection of different classes of organic compounds. Detection of minerals in vegetables or fruits means detection of elements other than carbon, hydrogen and oxygen.

MATERIAL REQUIRED

  • Test Tubes
  • Burner
  • Litmus paper
  • Laboratory reagents
  • Various fruits
  • Vegetables juices

CHEMICAL REQUIREMENTS

  • pH indicator
  • Iodine solution
  • Fehling solution A and Fehling solution B
  • Ammonium chloride solution
  • Ammonium hvdroxide
  • Ammonium oxalate
  • Potassium sulphocynaide solution

PROCEDURE

The juices are made dilute by adding distilled water to it, in order to remove colour and to make it colourless so that colour change can be easily watched and noted down. Now test for food components are taken down with the solution.

TEST, OBSERVATION & INFERENCE

Test
Observation
Inference
ORANGE TEST:
Test for acidity:
Take 5ml of orange juice in a test tube and dip a pH paper in it. If pH is less than 7 the juice is acidic else the juice is basic.The pH comes out to be 6.Orange juice is acidic.
Test for Startch:
Take 2 ml of juice in a test tube and add few drops of iodine solution. It turns blue black in colour than the starch is present.Absence of blue black in colour.Orange juice is acidic.
Test for Carbohydrates (FEHLING’S TEST):
Take 2 ml of juice and 1 ml of fehling solution A & B and boil it. Red precipitates indicates the presence of producing sugar like maltose, glucose , fructose & Lactose.No red coloured precipitates obtained.Carbohydrates absent.
Test for Iron:
Take 2 ml of juice add drop of conc. Nitric acid. Boil the solution cool and add 2-3 drops of potassium sulphocyanide solution .Blood red colours shows the presence of iron.Absence of blood red colour.Iron is absent.
Test for Calcium:
Take 2 ml of juice add Ammonium chloride and ammonium hydroxide solution. Filter the solution and to the filterate add 2 ml of Ammonium Oxalate solution. white ppt or milkiness indicates the presence of calcium.Yellow precipitate is obtained.Calcium is present.

CONCLUSION

From the table given behind it can be conducted that most of the fruits & vegetable contain carbohydrate & vegetable contain carbohydrate to a small extent. Proteins are present in small quantity. Therefore one must not only depend on fruits and vegetables for a balance diet

TO STUDY THE RATE OF EVAPORATION OF DIFFERENT LIQUIDS

INTRODUCTION

When a liquid is placed in an open vessel, it slowly escapes into gas phase, eventually leaving the vessel empty. This phenomenon is known as evaporation. Evaporation of liquids can be explained in terms of kinetic molecular model. Although there are strong inter-molecular attractive forces which hold molecules of a liquid together, the molecules having sufficient kinetic energy can escape into gas phase if such molecules happen to come near the surface. In a sample of liquid all the molecules do not have same kinetic energy. There is a small fraction of molecules which have enough kinetic energy to overcome the attractive forces and escape into gas phase.
Evaporation causes cooling. This is due to the reason that the molecules, which undergo evaporation, are high-energy molecules; therefore the kinetic energy of molecules which are left behind is less. Since the remaining molecules have lower average kinetic energy therefore, temperature must be lower. If the temperature is kept constant the remaining liquid will have the same distribution of molecular kinetic energies and the high-energy molecule will keep on escaping from the liquid into the gas phase. If the liquid is taken in an open vessel, evaporation will continue until whole of the liquid evaporates.

REQUIREMENTS

Apparatus:

  • Three petridishes of diameter 10 cm with covers
  • 10 ml pipette
  • Stop watch

Chemicals:

  • Acetone
  • Benzene
  • Chloroform

PROCEDURE

  • Clean and dry the petridishes and mark them as A, B, C.
  • Pipette out 10 ml of acetone to petridish A and cover it.
  • Pipette out 10 ml of benzene in petridish B and cover it.
  • Pipette out 10 ml of chloroform in petridish C and cover it.
  • Uncover all the three petridishes simultaneously and start the stop-watch.
  • Note the respective time when the liquids evaporate completely from each petridish.

OBSERVATIONS

Petridish Mark
Liquid Taken
Time taken for complete evaporation
AAcetone53 min
BBenzene42 min
CChloroform30 min

CONCLUSION

The rate of evaporation of the given three liquids is in the order:
Chloroform > Benzene > Acetone

Paper Chromatography: Basic Version

Objective
The objective of this project is to use paper chromatography to analyze ink components in permanent black markers.
Introduction
Matter makes up everything in the universe. Our body, the stars, computers, and coffee mugs are all made of matter. There are three different types of matter: solid, liquid, and gas. A solid is something that is normally hard (your bones, the floor under your feet, etc.), but it can also be powdery, like sugar or flour. Solids are substances that are rigid and have definite shapes. Liquids flow and assume the shape of their container; they are also difficult to compress (a powder can take the same shape as its container, but it is a collection of solids that are very small). Examples of liquids are milk, orange juice, water, and vegetable oil. Gases are around you all the time, but you may not be able to see them. The air we breathe is made up of a mixture of gases. The steam from boiling water is water’s gaseous form. Gases can occupy all the parts of a container (they expand to fill their containers), and they are easily compressed.
Matter is often a mixture of different substances. A heterogeneous mixture is when the mixture is made up of parts that are dissimilar (sand is a heterogeneous mixture). Homogeneous mixtures (also called solutions) are uniform in structure (milk is a homogeneous mixture). A sugar cube floating in water is a heterogeneous mixture, whereas sugar dissolved in water is a homogeneous mixture. You will determine whether the ink contained in a marker is a heterogeneous or homogeneous mixture, or just one compound.
In a mixture, the substance dissolved in another substance is called the solute. The substance doing the dissolving is called the solvent. If you dissolve sugar in water, the sugar is the solute and the water is the solvent.
For this project, you will be making a small spot with an ink marker onto a strip of paper. The bottom of this strip will then be placed in a dish of water, and the water will soak up into the paper.
The water (solvent) is the mobile phase of the chromatography system, whereas the paper is the stationary phase. These two phases are the basic principles of chromatography. Chromatography works by something called capillary action. The attraction of the water to the paper (adhesion force) is larger than the attraction of the water to itself (cohesion force), hence the water moves up the paper. The ink will also be attracted to the paper, to itself, and to the water differently, and thus a different component will move a different distance depending upon the strength of attraction to each of these objects. As an analogy, let’s pretend you are at a family reunion. You enjoy giving people hugs and talking with your relatives, but your cousin does not. As you make your way to the door to leave, you give a hug to every one of your relatives, and your cousin just says “bye.” So, your cousin will make it to the door more quickly than you will. You are more attracted to your relatives, just as some chemical samples may be more attracted to the paper than the solvent, and thus will not move up the solid phase as quickly. Your cousin is more attracted to the idea of leaving, which is like the solvent (the mobile phase).
Chromatography is used in many different industries and labs. The police and other investigators use chromatography to identify clues at a crime scene like blood, ink, or drugs. More accurate chromatography in combination with expensive equipment is used to make sure a food company’s processes are working correctly and they are creating the right product. This type of chromatography works the same way as regular chromatography, but a scanner system in conjunction with a computer can be used to identify the different chemicals and their amounts. Chemists use chromatography in labs to track the progress of a reaction. By looking at the sample spots on the chromatography plate, they can easily find out when the products start to form and when the reactants have been used up (i.e., when the reaction is complete). Chemists and biologists also use chromatography to identify the compounds present in a sample, such as plants.
Terms, Concepts and Questions to Start Background Research
  • adhesion, cohesion forces
  • capillary action
  • stationary phase, mobile phase
  • hydrophilic, hydrophobic
  • Rf value
  • paper chromatography
  • solvent
  • solution
Questions
  • Why do different compounds travel different distances on the piece of paper?
  • How is an Rf value useful?
  • What is chromatography used for?
Bibliography
Materials and Equipment
  • water
  • at least 15 identically sized strips of paper (5 for each pen)
    Note: chromatography paper or laboratory filter paper is preferable, but you can use a paper towel. The problem with paper towels is that they may be too absorptive and smear the ink. For more information on which papers work and which don’t.
  • ruler
  • pencils
  • at least three different types of black markers (including one permanent marker), or at least three different colors of marker (including one permanent marker)
  • a wide-mouth jar for the solvent
Experimental Procedure
Note: To make sure you can compare your results, as many of your materials as possible should remain constant. This means that the temperature, type of water used, size of paper strips, where the ink is placed onto the paper etc. should remain the same throughout the experiment.
  1. Cut paper strips about one by four inches in area (they must all be the same size).
  2. Take one of the paper strips and use a ruler and pencil to draw a line across it horizontally two cm from the bottom. This is the origin line (see illustration, below).
    Origin-Spot Diagram
  3. Pour a small amount of water into your glass (there should be barely enough for the paper strip to hang inside of the jar and just touch the water).
  4. Using one of the markers, place a small dot of ink onto the line (see illustration, above).
  5. Use the pencil to label the strip, so that you know which marker it represents.
  6. Tape the paper to a pencil and hang it into the jar of solvent so that the bottom edge is just barely touching (see illustration, below).
    Glass-Paper Example
  7. Let the water rise up the strip until it is almost at the top.
  8. Remove the strip from the jar and mark how far the solvent rose with a pencil.
  9. Analyze the ink component(s):
    Measure the distance the solvent and each ink component traveled from the starting position, then calculate the Rf value for each component (some of the ink components might not have moved at all!).
  10. Repeat this experiment for each brand or color of marker five times.
Questions
  • Did the different inks separate differently? By looking at the Rfvalues, can you tell if any of the ink components from the different markers are the same?
  • If the ink components separated differently for each marker, why did this happen (think about the strength of attractions)?

What’s the Point of Boiling?

Objective
The goal of this project is to separate pure water from fruit juice using a simple stovetop distillation apparatus.
Introduction
This project uses the technique of distillation. Distillation is when you boil a liquid, and then capture the vapor that escapes from the liquid and cool it. The cooled vapor condenses back into liquid. The condensed liquid is called the distillate. Do you think this process changes the liquid?
What if the liquid you boil has substances dissolved in it? For example, what if you started with a solution of sugar water? If you boiled the sugar water, you know from experience that there would be steam rising up from the pot on the stove. If you condensed that steam back into liquid, do you think the condensed liquid (the distillate) would contain sugar or not?
In this project, you will learn how to build a simple stove top distillation apparatus with stuff that you probably have in your kitchen right now. All you need is a deep pot with a sloping lid, a coffee cup, a bowl, some ice, and a stove. Of course, you’ll also need a liquid to distill. Colored fruit juice will work fine, or you could make a solution of sugar water. Add food coloring to it if you like. The Experimental Procedure section, below, shows you how to put it all together to find out what happens.
Terms, Concepts and Questions to Start Background Research
To do this project, you should do research that enables you to understand the following terms and concepts:
  • Boiling point
  • Phases of matter:
    • Solid
    • Liquid
    • Vapor
  • Condensation
  • Solvent
  • Solute
  • Distillate
Questions
  • What happens to solute molecules when the solvent evaporates or boils?
  • How will the distillate compare to the original juice for:
    • color?
    • taste?
    • pH?
Bibliography
Materials and Equipment
To do this experiment you will need the following materials and equipment:
  • Stove
  • Deep cooking pot with sloped lid
  • Ceramic coffee cup
  • Ceramic bowl
  • Ice
  • Hot mitts
  • Colored fruit juice (e.g., orange juice, grape juice, cranberry juice, etc.)
Experimental Procedure
  1. Do your background research so that you are familiar with the terms, concepts, and questions, above. For more information on distillation methods.
  2. The line drawing below is an illustration of the stove-top distillation apparatus used in this experiment.
    stove top distillation apparatus
    Line drawing of a stove top distillation apparatus. The text explains how to use it.
  3. Here are the steps for using the distillation apparatus.
    1. Pour the colored fruit juice into the bottom of the pot. Save at least 200 ml of the original juice for comparison to the distillate.
    2. Place the ceramic coffee cup, open side up, in the center of the deep pot. (That’s correct, right in the juice!)
    3. Place a bowl on top of the coffee cup. (The bowl will catch the condensed liquid that drips down from the lid.)
    4. Put the cover on the pot, upside down.
    5. Put ice in the cover of the pot.
    6. Turn on the burner to medium heat. You want the juice to boil moderately (not a rolling boil).
    7. Allow the pot to boil for 10 minutes or so (enough time to collect a sufficient amount of distillate for testing).
    8. When done, turn off the burner. Allow the pot to cool for a few minutes.
    9. Put on hot mitts and carefully remove the cover from the pot.
    10. Still wearing hot mitts, lift the bowl off of the coffee cup and set it down on a heat-resistant surface.
    11. Remove the coffee cup.
    12. After it cools, pour the remaining juice from the pot into a clear container.
  4. How do the original juice, the remaining juice from the pot, and the distillate compare in terms of color?
  5. Ordinarily in a chemistry experiment, you would not taste any of the solutions. In this case, since you are using clean kitchen utensils, and edible fruit juice, a taste test is OK. Let the liquids cool to room temperature before tasting them! How do the three different liquids compare for taste?
    1. Which liquid is sweetest?
    2. Which is least sweet?
    3. You should be able to explain why.

Make Your Own pH Paper

Objective
The goal of this project is to make your own pH indicator paper, and use it to measure the acidity and alkanity of various solutions from around your house.
Introduction
In this project you’ll learn how to make your own pH paper that you can use to find out if a solution is acidic or basic (alkaline). What does it mean for a solution to be acidic or alkaline?
It all has to do with hydrogen ions (abbreviated with the chemical symbol H+). In water (H2O), a small number of the molecules dissociate (split up). Some of the water molecules lose a hydrogen and become hydroxyl ions (OH). The “lost” hydrogen ions join up with water molecules to form hydronium ions (H3O+). For simplicity, hydronium ions are referred to as hydrogen ions H+. In pure water, there are an equal number of hydrogen ions and hydroxyl ions. The solution is neither acidic or basic.
An acid is a substance that donates hydrogen ions. Because of this, when an acid is dissolved in water, the balance between hydrogen ions and hydroxyl ions is shifted. Now there are more hydrogen ions than hydroxyl ions in the solution. This kind of solution is acidic.
A base is a substance that accepts hydrogen ions. When a base is dissolved in water, the balance between hydrogen ions and hydroxyl ions shifts the opposite way. Because the base “soaks up” hydrogen ions, the result is a solution with more hydroxyl ions than hydrogen ions. This kind of solution is alkaline.
Acidity and alkalinity are measured with a logarithmic scale called pH. Here’s why: A strongly acidic solution can have one hundred million million (100,000,000,000,000) times more hydrogen ions than a strongly basic solution! The flip side, of course, is that a strongly basic solution can have 100,000,000,000,000 times more hydroxide ions than a strongly acidic solution. Moreover, the hydrogen ion and hydroxide ion concentrations in everyday solutions can vary over that entire range. In order to deal with these large numbers more easily, scientists use a logarithmic scale, the pH scale. Each one-unit change in the pH scale corresponds to a ten-fold change in hydrogen ion concentration. The pH scale ranges from 0 to 14. It’s a lot easier to use a logarithmic scale instead of always having to write down all those zeros! By the way, notice how one hundred million million is a one with fourteen zeros after it? It’s not coincidence, it’s logarithms!
To be more precise, pH is the negative logarithm of the hydrogen ion concentration:
pH = log 1/[H]+ = −log [H+] .
The square brackets around the H+ automatically mean “concentration” to a chemist. What the equation means is just what we said before: for each 1-unit change in pH, the hydrogen ion concentration changes ten-fold. Pure water has a neutral pH of 7. pH values lower than 7 are acidic, and pH values higher than 7 are alkaline (basic). The table below has examples of substances with different pH values (Decelles, 2002; Environment Canada, 2002; EPA, date unknown).
Table 1. The pH Scale: Some Examples
pH ValueH+ Concentration
Relative to Pure Water
Example
010 000 000battery acid
11 000 000sulfuric acid
2100 000lemon juice, vinegar
310 000orange juice, soda
41 000tomato juice, acid rain
5100black coffee, bananas
610urine, milk
71pure water
80.1sea water, eggs
90.01baking soda
100.001Great Salt Lake, milk of magnesia
110.000 1ammonia solution
120.000 01soapy water
130.000 001bleach, oven cleaner
140.000 000 1liquid drain cleaner
In this project you will make your own pH paper from a colored indicator that you will extract from red cabbage by cooking it in water. Once you have the indicator solution, you can soak some coffee filter paper in it, then allow the paper to dry. When the paper is dry, you can cut it into strips, and you’ll have pH paper that will change color. It will turn greenish when exposed to bases, and reddish when exposed to acids. How green or how red? That’s your job! Use different solutions that you have around the house to find out how the color change corresponds to changes in pH.
Terms, Concepts and Questions to Start Background Research
To do this project, you should do research that enables you to understand the following terms and concepts:
  • Acids
  • Bases
  • Logarithms
  • pH
  • pH indicators
Questions
  • What value of pH is neutral?
  • What range of pH values is acidic?
  • What range of pH values is basic?
  • What color is red cabbage pH paper when dipped in acidic solutions?
  • What color is red cabbage pH paper when dipped in basic solutions?
Bibliography
Materials and Equipment
To do this experiment you will need the following materials and equipment:
  • Red cabbage leaves
  • 1-quart cooking pot
  • Water
  • 1-quart bowl
  • Strainer
  • White coffee filters (cone-shaped ones are good)
    • Alternatively, you can use filter paper or chromatography paper.
  • Acidic and basic solutions to test, for example:
    • Lemon juice, vinegar
    • Orange juice, soda
    • Tomato juice, acid rain
    • Black coffee, bananas
    • Milk, saliva
    • Pure water
    • Sea water, eggs
    • Baking soda solution
    • Milk of magnesia
    • Ammonia solution
    • Soapy water
Experimental Procedure
Safety Notes:
  • Adult supervision required.
  • Do not mix strong acids and bases.
  • Use appropriate caution when testing the pH of household cleaning solutions (like ammonia). Avoid skin contact, and follow all precautions on the product label.
  1. Do your background research so that you are knowledgeable about the terms, concepts, and questions, above.
  2. Prepare a red cabbage indicator solution (the “Experiments with Acids and Bases” webpage (Carboni, 2004) has great pictures illustrating all of the steps)
    1. Slice a head of cabbage at approximately 3 cm (1 in) intervals, or peel the leaves from the head and tear them into pieces.
    2. Place the leaves in the cooking pot and cover with water.
    3. Cook on medium heat for half an hour (low boil is good).
    4. Allow the cooked cabbage to cool, then pour off the liquid into a bowl. You can pour through a strainer to catch the cabbage pieces, or hold them back with a large, flat ladle with holes—see the photographs on the “Experiments with Acids and Bases” webpage (Carboni, 2004).
    5. The solution is a deep blue, but will change color when the pH changes. (You can experiment with using the liquid as a pH indicator.)
  3. Here’s how to make pH paper using the red cabbage solution and coffee filters:
    1. Soak the white coffee filters in the red cabbage solution for about 30 minutes.
    2. Drain the excess solution from the filters, and set them out in a single layer on some paper towels to dry overnight. To speed up the drying process, you can put them on a cookie sheet and put them in your oven at low temperature (150–200°F.
    3. When the coffee filters are dry, cut them into 3 cm × 8 cm (about 1 in × 3 in) strips.
    4. The strips are now ready to test the pH of various solutions. They start out blue, but will turn green in basic solutions and red in acidic solutions.
  4. Use the strips to test the acidity/alkalinity of various solutions around your house. For example:
    • Lemon juice, vinegar
    • Orange juice, soda
    • Tomato juice, acid rain
    • Black coffee, bananas
    • Milk, saliva
    • Pure water
    • Sea water, eggs
    • Baking soda solution
    • Milk of magnesia
    • Ammonia solution
    • Soapy water
    • Note: if you test the pH of saliva, do not put the pH paper in your mouth! Instead, spit some saliva into a clean container and dip the paper into the saliva.
  5. After testing, put the pH strips in order of increasing pH of the solution tested.
    1. You can use the table in the Introduction as a guide.
    2. The Variations section has some additional suggestions for independent confirmations of the pH readings.
  6. Do you see a gradual change in color as the pH of the tested solutions varies? Can you match specific colors to certain pH levels? Over what range of pH does the color continue to change? How accurately do you think you can determine the pH of a solution with your test papers? Within 1, 2, or 3 pH units?
Variations
  • Compare the performance of your homemade pH paper with commercial pH paper (can be found in a well-stocked tropical fish store). Or, buy an inexpensive pH meter and use it to calibrate your homemade pH paper. Use the table in the Introduction to make a series of different solutions, form low to high pH. Measure the pH of each solution with the pH meter (rinse off the tip between solutions), and write down the results. Now check each solution with your pH paper. Can you see color differences that correspond to the measured changes in pH? Over what pH range do you see color changes? How large does the shift in pH need to be in order to see a change in color?
  • Try making pH indicator solutions (and/or indicator papers) from other natural dyes: for example beet juice, phenolphthalein, or turmeric powder (Beckham, date unknown; Krampf, 2006). Test your household solutions with each of the indicators. Does the additional information from multiple indicators give you a better measure of the pH of your solutions?
  • Does the pH of your saliva change after eating various types of food? If so, how much time does it take to return to normal? Design an experiment to find out. Again, do not put the pH paper in your mouth. Instead, spit some saliva into a clean container and dip the pH paper into the saliva. Also, don’t try changing the pH of your saliva with anything non-edible!
  • What is the pH of rainwater in your area? Can you measure it with your pH paper or pH indicator solutions?

What Makes Ice Melt Fastest?

Objective
The goal of this project is to determine which added material will make ice melt fastest.
Introduction
To make ice cream with an old-fashioned hand-crank machine, you need ice and rock salt to make the cream mixture cold enough to freeze. If you live in a cold climate, you’ve seen the trucks that salt and sand the streets after a snowfall to prevent ice from building up on the roads. In both of these instances, salt is acting to lower the freezing point of water.
For the ice cream maker, because the rock salt lowers the freezing point of the ice, the temperature of the ice/rock salt mixture can go below the normal freezing point of water. This makes it possible to freeze the ice cream mixture in the inner container of the ice cream machine. For the salt spread on streets in wintertime, the lowered freezing point means that snow and ice can melt even when the weather is below the normal freezing point of water. Both the ice cream maker and road salt are examples of freezing point depression.
Salt water is an example of a chemical solution. In a solution, there is a solvent (the water in this example), and a solute (the salt in this example). A molecule of the solute will dissolve (go into solution) when the force of attraction between solute molecule and the solvent molecules is greater than the force of attraction between the molecules of the solute. Water (H2O) is a good solvent because it is partially polarized. The hydrogen ends of the water molecule have a partial positive charge, and the oxygen end of the molecule has a partial negative charge. This is because the oxygen atom holds on more tightly to the electrons it shares with the hydrogen atoms. The partial charges make it possible for water molecules to arrange themselves around charged atoms (ions) in solution, like the sodium (Na+) and chloride (Cl) ions that dissociate when table salt dissolves in water.
Other substances that dissolve in water also lower the freezing point of the solution. The amount by which the freezing point is lowered depends only on the number of molecules dissolved, not on their chemical nature. This is an example of a colligative property. In this project, you’ll investigate different substances to see how they affect the rate at which ice cubes melt. You’ll test substances that dissolve in water (i.e., soluble substances), like salt and sugar, as well as substances that don’t dissolve in water (i.e., insoluble substances), like sand and pepper. Which substances will speed up the melting of the ice?
Terms, Concepts and Questions to Start Background Research
To do this project, you should do research that enables you to understand the following terms and concepts:
  • Solution
  • Solute
  • Solvent
  • Colligative properties
  • Freezing point depression
  • Phases of matter
    • Solid
    • Liquid
    • Gas
    • Plasma
  • Phase transitions
    • Melting
    • Freezing
    • Evaporation
    • Condensation
    • Sublimation
Questions
  • Which of the suggested test substances are soluble in water?
  • Which of the suggested test substances are insoluble in water?
Bibliography
Materials and Equipment
To do this experiment you will need the following materials and equipment:
  • Ice cubes
  • Identical plates or saucers
  • Timer
  • Electronic kitchen balance (accurate to 0.1 g)
  • Measuring cup
  • Suggested materials to test for ice-melting ability
    • Table salt
    • Sugar
    • Sand
    • Pepper
Experimental Procedure
  1. Do your background research so that you are knowledgeable about the terms, concepts, and questions, above.
  2. You’ll need a clean plate and several ice cubes for each of the substances to be tested.
  3. Note the starting time, then carefully sprinkle one teaspoon of the substance to be tested over the ice cube.
  4. After a fixed amount of time (say, 10 minutes), pour off the melted water into a measuring cup, and use the balance to measure the mass. Subtract the mass of the empty cup, and you’ll have the mass of the melted water. Wait the same amount of time for each test.
  5. Measure the remaining mass of the ice cube.
  6. Repeat three times for each substance to be tested.
  7. Use the same procedure to measure the melting rate for ice cubes with nothing added.
  8. For each test, calculate the percentage of the ice cube that melted:[mass of melt water]/[initial mass of ice cube] × 100
  9. For each test, calculate the percentage of the ice cube remaining:[remaining mass of ice cube]/[initial mass of ice cube] × 100
  10. For each substance you tested, calculate the average amount of melted water produced (as a percentage of initial mass), and the average remaining ice cube mass (as a percentage of initial mass).
  11. Did any substances speed up melting of the ice (compared to melting rate of plain ice cubes with nothing added)?
Variations
  • Does the melting rate depend on the amount of solute added? Design an experiment to find out.
  • Try your experiment in the refrigerator to simulate colder weather. Alternatively, if the outside temperature is wintry, take your experiment outdoors! Be sure to monitor the temperature regularly throughout your experiment.
  • Do you think salt would melt ice in your freezer? Why or why not? Try it and find out.

Measuring Surface Tension of Water with a Penny

Objective
The goal of this project is to investigate how added salt and added detergent affect the surface tension of water.
Introduction
Water molecules—good old H2O—are made of one oxygen and two hydrogen atoms. The single oxygen and two hydrogen atoms are held together because they share electrons—this is called a covalent bond. The hydrogen atoms don’t line up on opposite sides of the oxygen atom, as you might think. Instead they are at an angle of about 105° (if they were on opposite sides of the oxygen atom the angle would, of course, be 180°).
The oxygen atom tends to hold on to the shared electrons from the hydrogen atoms more tightly, so each end of the water molecule ends up with a partial charge. The oxygen portion of the molecule has a partial negative charge, and the hydrogen ends of the molecule have a partial positive charge. Another way of talking about the partial charges is to say that water molecules are polarized. Like a magnet, with a north and south pole, a water molecule has electrical poles. The oxygen atom is the negative pole, and each hydrogen atom is a positive pole.
These partial charges cause water molecules to interact with one another. Because opposite charges attract, water molecules tend to ‘stick’ to one another. The partial positive charges of the hydrogen atoms tend to align themselves with the partial negative charge of the oxygen atoms of neighboring water molecules. You can see models of this alignment in several of the references in the Bibliography section (Wiseth, date unknown; Hipschman, 1995a; Kimball, 2006). This tendency of water molecules to stick together due to the partial positive and negative charges is called hydrogen bonding.
Hydrogen bonding between water molecules leads to many interesting consequences at the visible, macroscopic level. For example: the boiling point of water, its surface tension, and it’s ability to dissolve salts are all related to hydrogen bonding.
The boiling point of water, 100°C, is unusually high for a molecule with such a low molecular weight. The boiling point is so high due to hydrogen bonding. On average, each water molecule interacts with about four others (each hydrogen atom interacts with the oxygen atom of separate water molecules, and each oxygen atom interacts with the hydrogen atoms of two more water molecules). In water vapor, the molecules are too far apart for hydrogen bonding to occur, so boiling water means breaking up all of the hydrogen bonds in liquid water. Breaking those bonds takes energy, thus the high boiling point for water.
Hydrogen bonds also give liquid water a high surface tension. The water molecules on the surface have partners for hydrogen bonding only within the liquid; above the water surface there are no more molecules available for hydrogen bonding. This means that molecules at the surface experience a net force pulling them inward. If you fill a glass right up to the rim and then carefully add a few more drops of water, you can see that the glass can be overfilled without spilling. The surface tension of the water holds on to the ‘extra’ water as if there were a skin on the surface of the water.
Water is an excellent solvent for charged (polar) molecules like table salt, NaCl. In water, salt dissociates into positively charged sodium (Na+) and negatively charged chloride (Cl) ions. The partial positive charge of the hydrogen ends of the water molecules surround the negatively charged chloride ions, and the partial negative charge of the oxygen ends of the water molecules surround the positively charged sodium ions. What effect will dissolved salt ions have on hydrogen bonds between water molecules?
Water behaves very differently when mixed with uncharged (nonpolar) molecules. An example of a nonpolar molecule is cooking oil. You may have heard the saying “oil and water don’t mix,” and this is why. Oil molecules are uncharged. Water molecules, as you have learned, are partially charged. The uncharged oil molecules disrupt the hydrogen bonding between water molecules. So when you try to mix oil and water, the oil ends up forming droplets within the water. The nonpolar oil molecules stick together and the polar water molecules stick together. Eventually, you get two layers, with the less dense oil floating on top of the denser water.
Nonpolar substances are sometimes called ‘hydrophobic’ (meaning ‘water fearing’), and polar molecules are sometimes called ‘hydrophilic’ (meaning ‘water loving’) because of the two different interactions illustrated by salt and cooking oil.
Liquid detergents have dual properties. One end of the molecule is oily, and the other end is charged. In water, the oily ends of detergent molecules stick together, with the charged ends sticking out, into the water. Detergents can form small blobs in water (called micelles) and can also disperse, like oils, into a layer on the surface of the water (for illustrations, see Hipschman, 1995b). How do you think added detergent will affect the surface tension of water?
One way to find out is to count how many drops of water you can ‘pile up’ on top of a single penny. The Experimental Procedure section shows you how to do this with plain water, salt water, and water with detergent.
Terms, Concepts and Questions to Start Background Research
To do this project, you should do research that enables you to understand the following terms and concepts:
  • surface tension,
  • chemical structure of water,
  • covalent bond,
  • hydrogen bonds,
  • polar solvent,
  • non-polar solvent,
  • hydrophobic,
  • hydrophilic.
Questions
  • What happens to salt when it is dissolved in water?
  • How do you think adding salt to the water will affect the hydrogen bonds between water molecules?
    • What effect will this have on surface tension?
    • Do you think salt water will have more or less surface tension than plain tap water?
  • What happens to detergent when it is dissolved in water?
  • How will added detergent affect hydrogen bonds between water molecules?
    • What effect will this have on surface tension?
    • Do you think water with added detergent will have more or less surface tension than plain tap water?
Bibliography
Materials and Equipment
To do this experiment you will need the following materials and equipment:
  • water,
  • plastic transfer pipettes (or eyedropper),
    • one online source of transfer pipettes in small quantities is RachelsSupply.com, where you can get 10 fine-tipped pipettes (part #1N01) for $1.50 + shipping.
  • salt,
  • dishwashing detergent,
  • clean glass jars (or beakers),
  • measuring spoons.
Experimental Procedure
  1. Holding the transfer pipette close to the surface of the penny, carefully pipet water droplets onto the penny, one at a time, counting each drop. Tips:
    1. The droplets should pool up on the penny, creating a big droplet of water.
    2. To make sure your count is accurate, hold the pipette far enough above the penny so that the drop has to fall a short distance before fusing with the droplet on the penny.
  2. Stop pipetting when the droplet on the penny breaks up and overflows. The count for each trial is the number of drops that the penny could hold (in other words, count all of the dropsexcept the one that caused the penny to overflow).
  3. Repeat the measurement ten times for each solution that you test.
  4. Test the following solutions:
    1. added salt: dissolve 1 teaspoon (6 grams) in 100 mL of water,
    2. added detergent: put 1 drop of liquid dishwashing detergent in 1 liter of water; do not shake–cap the container and gently tip it back and forth to mix.
Variations
  • Try a series of increasing concentrations of salt (maximum solubility at room temperature is about 36 g salt/100 mL water). The best way to do this is by making a concentrated solution, and then making serial dilutions to make less-concentrated solutions. Does surface tension continue to change as more salt is added? Students who have studied high school chemistry should compare molar ratios of NaCl and H2O.
  • Do you think that changing the temperature of the water would affect surface tension? How? Design an experiment to find out. Measuring surface tension on a penny is probably not the best design for this variation, because the temperature would not be well controlled. The volume of water is quite small, so the temperature could easily change. However, if you controlled the temperature of the water and the penny, you could probably get this to work. Another idea would be to find a different way to measure surface tension, using a larger volume of water.
  • Does the surface of the penny matter? What happens if you coat the penny with a thin film of cooking oil? Wet a paper towel slightly with cooking oil. Wipe off the excess oil, then use the paper towel to wipe a thin film of oil on a penny. How many drops of water will the penny hold compared to a “normal” penny. Can you think of other surface treatments you could try? Could you make penny-sized disks of other materials to test? How important is the raised edge of the penny for holding the water? Does the ‘heads’ side hold more or less water than the ‘tails’ side?