Lab Activity No. 2

Chemistry and Water

General Biology Lab

A. Objectives: Upon completion of this lab activity, you should be able to:

1.Explain the different types of chemical bonds (covalent, polar covalent, ionic, and hydrogen bonds).

2.Explain how different chemicals can interact with one another resulting in an alteration of the properties of the chemical system.

3.Describe the properties of water that result from the hydrogen-bonding ability of water molecules.

B. Introduction: Chemistry, the interactions and reactions of atoms to form molecules, is the basis for all of life. Living organisms are essentially bags of chemicals--specific chemicals that perform all of the functions necessary to keep their collective body alive. In this lab, we will examine fundamental properties of chemicals. We will consider the different types of chemical bonds that can form and how these bonds affect the larger properties of chemicals, such as water, with which we are more familiar.

C. Molecular Movement and Chemical Reactions (Demonstration)

Molecules are in constant motion. Solid materials are made of molecules that are bound together in such a way that they move slowly and not very far. Liquids are made of molecules that are bound together, but not very strongly. Gases are molecules that are independent and able to move rapidly and over long distances. The lighter the molecule, the faster and farther it can travel. If two molecules bump into each other, they may react to form new molecules (products) that are more stable (less reactive).

Materials:

  • One glass tube, held horizontally in a stand
  • Cotton plugs
  • Concentrated HCl (Hydrochloric Acid)
  • Concentrated NH4OH (Ammonium Hydroxide)

1.Place cotton plugs in both ends of the glass tube.

2.Drench one cotton plug with NH4OH. Drench the other cotton plug with HCl.

3.Allow the system to react for a few minutes.

Describe the results. What products formed (describe what you see)? Where did they form (right in the middle or off to one side or the other)?

Why did these things happen? Explain using the terms molecular movement and chemical reaction.

D. The Polar Nature of Water

Water is an amazing substance. Water molecules are made of two hydrogen atoms and one, oxygen atom. These atoms form two polar covalent O-H bonds. Water is therefore a polar molecule with a negatively charged end and a positively charged end. Water molecules attract one another--the negative end of one molecule attracts the positive end of another water molecule. These attractions are called hydrogen bonds and they are what give water its unique properties.

Materials:

  • A glass rod
  • Fur or wool
  • A buret
  • A ring stand with clamps
  • 400 ml beaker

1. Fill a buret with water and place it in the clamp of the ring stand with a 400 ml beaker underneath.

2. Vigorously rub a glass rod with fur or wool for 1 minute.

3. Open the buret so that a very gentle, fine stream of water is produced.

4. Place the prepared rod about 1 cm from the stream and observe the stream carefully. Be careful not to touch the water stream.

5. Now place the rod on the other side of the stream. What did you observe?

Draw what must be happening to the water at the molecular level. Hint: You are making the glass rod negatively charged by adding electrons to it as you rubbed it.

E. Adhesion

Adhesion occurs when water molecules form hydrogen bonds with other types of molecules.

Materials:

  • 2 microscope slides

1. Obtain two microscope slides and wash them until they are “squeaky clean” and dry them.

2. Place a single drop of water on one slide and sandwich it with the other slide.

3. Try to pull the two slides straight apart without sliding them against each other. The two slides are held together by the hydrogen bonding of water to glass. That’s adhesion!

F. Cohesion and Surface Tension

Cohesion results when water molecules form hydrogen bonds with other water molecules. Cohesioncreates surface tension on the surface of a body of water.

Materials:

  • 1 petri dish
  • 1 piece of lens paper
  • 1 straight pin
  • 2 toothpicks
  • Detergent in a dropper bottle

1. Fill the bottom of a Petri dish with water.

2. Float a small piece of lens paper (about 1cm x 4cm) on the water and then carefully place a straight pin on the lens paper.

3. Gently submerge the lens paper using two toothpicks. What did you observe? Explain what happened in terms of hydrogen bonding.

4. Add a drop of detergent as far away from the pin as possible. What did you observe? Explain what happened.

G. Specific Heat of Water

Heat is a result of molecular motion. The faster the molecules are moving, the more heat a substance contains. Temperature is a measurement of the amount of heat (how hot or cold something is). Water is able to maintain a relatively stable temperature because the thermal motion of the molecules is slowed by hydrogen bonds between the molecules. Therefore, it takes a considerable amount of heat to raise the temperature of water. For a given amount of heat input, the temperature of water will rise more slowly than for virtually any other substance. Conversely, the temperature will fall more slowly once heat is removed because much of the energy lost is being used to reform hydrogen bonds between molecules, which is not recorded as a temperature decrease on the thermometer. Thus it is said that water has a high specific heat (the amount of energy required to raise or lower the temperature of a substance by 1C). The high specific heat of water stabilizes its temperature. We can observe this by comparing the cooling rate of water and oil.

Materials:

  • Colored mineral oil
  • 2, 125 ml Erlenmeyer flasks
  • 1 hot plate
  • Vented stoppers with thermometers
  • Boiling chips
  • 2, 400 ml beakers

1. Fill a 125 ml Erlenmeyer flask up to the neck with water. Fill a second flask with colored mineral oil. Fit the flasks with the rubber stoppers and thermometers. The thermometers should extend about two-thirds of the way into the water or oil.

2. Add 200 ml of water and a few boiling chips to each of two 400 ml beakers.

3. Place each of the Erlenmeyer flasks prepared in step 1 into the 400 ml beaker. The water level should be near the neck of the flask.

4. Heat the 400 ml beakers on hot plates until the temperature of the water or oil reaches 75 C. Then remove the beaker from the hot plate and take the flask out of the water. (CAUTION: Hot glassware looks just like cold glassware so when in doubt, use an oven mitt to avoid being burned!) Take the first temperature reading at 75 C and begin timing. Record the temperature at two-minute intervals for 20 minutes. Repeat this procedure for the other flask once it reaches 75 C. The data for the oil and water DO NOT need to be recorded simultaneously.

Data Table:

Time (min.) / Water Temp. (C) / Oil Temp. (C)
0
2
4
6
8
10
12
14
16
18
20

Graph of Results:

H. Temperature and Density (Demonstration)

Temperature affects the density of water. Your instructor will do the following demonstration to investigate this. Observe the results and answer the questions below.

Materials:

  • Water with ice cooled to between 1C and 4 dyed yellow
  • Warm water between 30C and 35C dyed blue

1. Pour 40 ml of the cooled, yellow water into a 100 ml graduated cylinder.

2. Gently pour 40 ml of the warm, blue water on top of the cold water. It will help to tip the graduated cylinder at an angle and pour the warm water down the side of the cylinder.

Describe what happened. Which color of water has the greater density? How does temperature affect the density of water in terms of molecular motion?

What happens to the density of water when it is cooled to 0C? What has happened on a molecular level? What bonds were involved? Hint: Does ice sink or float?

I. Acids and Bases (pH)

Water, as stated above, is made of two hydrogen atoms and one oxygen atom. In a glass of water, there are billions and billions of water molecules. As they move around in the glass, the water molecules push and shove one another. Sometimes a water molecule breaks into an H+ ion and an OH- ion. The H+ ion is missing an electron and the OH- ion has an extra electron. In a glass of pure water, at any moment in time, approximately 1 in 10 million (10-7) water molecules is broken into ions. These ions can react with one another to re-form a water molecule. So water molecules are continually breaking up and getting back together.

Under some circumstances, there are more H+ ions than OH- ions (or vice versa). In these situations, the solution is called either an acid (if there is an excess of H+ ions) or an alkaline or base (if there is an excess of OH- ions). The pH scale is used to measure the amount of H+ ions in a solution. The pH of pure water is 7 (the exponent in the scientific notation in the previous paragraph). Acids have a pH lower than 7; alkaline solutions have a pH greater than 7.

Materials:

  • pH test paper

1.Test the pH of each of the following chemicals. Record the results.

pH / Which is in excess, H+ or OH- or is H+=OH-?
  1. dH2O

  1. dH2O + CO2 (breath into water through a straw)

  1. Human saliva

  1. Vinegar

  1. Ammonia

  1. Lemon juice

  1. Apple juice

  1. Liquid Plummer

J. The Reactivity of Acids

As discussed above, acids have an excess of H+ ions. These ions, with their full positive charge, are very reactive. They are attractive to negatively charged ions and to the negative side of polar covalent bonds.

The three main types of organic compounds that make up living organisms are proteins, lipids (fats and oils), and carbohydrates (sugars and starches). These organic compounds are all built from many atoms that are covalently bound to one another.

Materials:

  • 50 ml beaker
  • Vinegar
  • 10% Lactose solution (carbohydrate)
  • Cream (lipid)
  • Nonfat milk (carbohydrate + protein)

1.Pour some lactose solution into a beaker. Add an equal amount of vinegar. Swirl to mix the two liquids. Describe the results below.

2.Rinse out the beaker and repeat step 1 using cream and then again using nonfat milk (adding vinegar to each).

Describe the results:

Lactose SolutionCreamNonfat Milk

Explain the results by describing the presence or absence of negative charges associated with each of the three types of organic compounds. (Which organic compound(s) contain(s) negative charges?)

What is the role of stomach acid (hydrochloric acid, pH = 2)?

Why does milk curdle when it goes bad? What causes milk to go bad and why does the milk solidify?

K. Instant Freeze (This activity was adapted from Bare, W. D. J. Chem Ed. 1991, 68, 1038 and from Flinn Scientific, Inc.)

When water is cooled the water molecules slow down and become arranged such that each molecule forms as many hydrogen bonds with other water molecules as possible. This results in a very regular arrangement of molecules; we recognize the orderly arrangement as ice crystals. Pure water freezes at 0C.

Dissolved substances (solutes) disrupt the hydrogen bonds of water molecules. In order for hydrogen bonds to form, the water molecules have to be slowed down even more than usual. Thus, if there are solutes present, the freezing point is lower than 0C.

Materials:

  • Club soda, 10 oz. in a clear glass bottle, cooled to refrigerator temperature
  • Thermometer, -20 to 110 °C
  • Beaker, 1-L
  • Crushed ice
  • Rock salt

1.Place a thin layer of crushed ice at the bottom of a 1-liter beaker.

2. Sprinkle a thin layer of rock salt into the beaker, overlaying the ice.

3. Place the cooled bottle of club soda in the center of the beaker and continue alternating layers of ice and rock salt around the bottle. Be sure to completely cover the bottle.

4. Place a thermometer into the ice/salt mixture. The thermometer should be very close to or touching the bottle in order to obtain an accurate reading of the temperature of the mixture affecting the soda.

5. The club soda must reach a temperature of -8 °C (17.6 °F) and remain there for about 10 minutes. Do not allow the soda to get too cold. The soda may freeze, ruining the demonstration. If the soda gets too cold (<-10 °C), there is a possibility of the bottle exploding.

6. After 10 minutes, observe that the soda is still a liquid at this point even though pure water would have frozen at this temperature.

7. Open the lid on the bottle while keeping the bottle in the ice. Notice how the club soda quickly solidifies.

Why did the temperature of the ice (and presumably the club soda too) drop below 0 °C? Ice should be at a temperature of 0 °C).

Why did the club soda remain liquid even at sub-zero temperatures?

Why did the club soda freeze solid almost immediately after the lid was popped open?

The following activity is optional and can be tried at home.

L. Plastic Bag Ice Cream (This activity was developed by Kimberly Granatire and Phillip Murry at a ICE Workshop at Miami University, Middletown, Ohio in July 1991.)

During the winter, in some parts of the country, icy roads are treated with salt to melt the ice.

In order for salted ice to melt, the water molecules must absorb heat from their surroundings. Thus, although salted ice melts (gets warmer) the surroundings (like the road or the concrete of your driveway) actually get colder. This is the principle that underlies ice cream making.

Materials:

  • 1 small zip lock plastic bag
  • 1 large zip lock plastic bag
  • 1/4 cup sugar
  • 1/2 cup (120 ml) milk
  • 1/2 cup (120 ml) heavy whipping cream
  • 1/4 teaspoon vanilla
  • 1 plastic spoon
  • 1/2 to 3/4 of a cup of rock salt
  • 3 to 4 cups of crushed ice
  1. Put sugar, milk, whipping cream, and vanilla into the small plastic bag. Carefully transfer the contents of the cup into the small zip lock bag. Seal the bag tightly so there are no leaks.
  2. Place the smaller plastic bag inside the larger bag.
  3. Surround the smaller bag with several cups of crushed ice.
  4. Pour 1/2 to 3/4 of a cup of salt over the ice and seal the larger bag securely.
  5. Knead or roll back and forth on a table or desktop. Be careful not to put too much pressure on the bags.
  6. After 10 minutes check the mixture to see if it is frozen. If not, continue kneading.
  7. When the mixture is frozen, simply remove the smaller bag and eat the ice cream directly from the bag. (Add nuts, fruit or chocolate if desired.)

Taste Good?

Why was it necessary to add rock salt to the ice? What would have resulted if the salt had not been added?

M. Review Questions

  1. Describe the general structure of an atom in terms of the relative locations of the three subatomic particles.
  1. What do the atomic number and atomic mass tell you about the atomic structure?
  1. Arrange the following substances in order from least amount of atomic motion to most: gas, solid, liquid.
  1. Write the chemical formula for both water and table salt. What 2 pieces of information are given to you in these formulas?
  1. What is the difference between an ionic bond, a non-polar covalent bond, a polar covalent bond, and a hydrogen bond? Which are strongest and which are weakest?
  1. What types of bonds occur in pure water? In pure salt? In salt water?
  1. Explain why salt dissolves in water.
  1. What do the terms “acid” and “base” mean? How does the pH scale measure this?
  1. List some properties of water that make it conducive to life. Give some specific examples.
  1. Why does water have a high specific heat? What is it about the molecules that cause this?

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