The Chemistry of Life

All life exists within the context of its environment. Each environment is characterized by its biological, physical, and chemical properties. Since organisms are adapted to a specific environment, radical changes in these conditions often result in injury to the individual or possibly extinction of the species. Recent reports of declining frog populations, for example, have been correlated with increased ultraviolet radiation from the sun (specifically UVB). Chemical reactions that take place inside of an organism are dependent upon both internal and external chemical and physical properties. We will explore some of these properties in today’s lab.

Part 1: pH Chemistry

Although water is generally regarded as a stable compound, individual water molecules are constantly gaining, losing, and swapping hydrogen atoms. This process is represented by the following chemical reaction:

H2O ↔ OH + H

Pure water with a pH of 7 has equal numbers of hydrogen and hydroxide ions at any given moment. Water is considered to be a neutral substance. The pH of any solution can be determined by calculating the total concentration of hydrogen ions in the solution.

What is a MOLE?

A mole is a term used to describe the quantity of something. If you have a mole of paperclips, that means you have paperclips. This is similar to the way we use the word “dozen.” We know that if you have a dozen paperclips, it means you have 12 of them.

1 dozen = 12

1 mole = 6.02×1023

Scientists measure acidity using the pH scale. The pH scale ranges from 0 to 14, and the numbers represent the concentration of hydrogen ions–1 in the substance. For example, battery acid, with a pH of 1, has 1×10 moles of hydrogen ions per liter of solution. Ammonia, which is a very basic substance with a pH of 12, has 1×10–12 moles of hydrogen ions per liter of solution. The more acidic the solution, the more hydrogen ions it contains.

Excessive changes in pH can cause metabolic and ecological problems. For example, the pH of your blood is carefully kept between 7.35 and 7.45. Any deviation above or below this range will result in alkalosis or acidosis, and both conditions can be deadly. Acid rain, on the other hand, can dissolve toxic metals from the soil particles into the soil solution and impair plant growth. As we will soon see, plant health is a factor that quickly affects most other life forms on the planet.

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  • pH paper (1–14)
  • Plastic tray with wells
  • Wax pencil


Following the instructions given by your teacher, measure the pH of each solution using pH paper.  Record the pH of these items below.

  • Tap water
  • Distilled water
  • Bleach
  • Aspirin
  • Lemon juice
  • Milk

Part 2: Buffers

Buffers are molecules that resist changes in pH. They can take up and release excess hydrogen ions in a solution and therefore prevent drastic changes in pH, regardless of whether acid or base is added to the solution. The net result is that the pH of the solution remains relatively stable (until the buffer is overwhelmed). Buffers are commonly found in dissolved minerals, soils, and in living organisms.

For example, buffers can play a protective role in lake ecosystems. In a lowland lake, acid rain causes very little fluctuation in pH because these lakes are typically high in organic molecules that act as buffers. A lake with little buffering capacity, such as a high alpine lake low in organic molecules, will experience a much greater change in overall pH as a result of acid precipitation.


  • pH paper (1–14)
  • Tap water (H2O)
  • 2 beakers (250 mL)
  • 1g NaHCO3 (baking soda)
  • 0.001 M HCl (hydrochloric acid)
  • Wax pencil


  1. Fill two beakers with 50 ml of tap water. Label one beaker “buffered” and label the other beaker “unbuffered.”
  2. Add 1 gram of baking soda to the “buffered” beaker. Swirl to dissolve.
  3. Using the pH strips, measure the pH of both beakers. Record all measurements in Table 1.
  4. Add 10 ml of 0.001 M HCl (hydrochloric acid) to each beaker and swirl.
  5. Measure the pH of the two beakers and record.
  6. Repeat steps 3–5 until you have added a total of 50 ml of 0.001 M HCl to each beaker.


Record the pH of your buffered and unbuffered solutions after each addition of 10 mL of hydrochloric acid to each beaker.

Table 1. Buffered and Unbuffered Solutions
Volume (mL) of
0.001 M HCl
Buffered Solution Unbuffered Solution

Data Analysis

Illustrate the buffering capacity of each solution by graphing your results below. Place the volume of  HCl on the x-axis and the pH value on the y-axis. Don’t forget to give your graph a title.

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You can download this graph paper template to complete this portion.

Part 3: Buffers in the Blood

Bicarbonate ions act as a powerful buffer in your blood. They are created when carbon dioxide (CO2), produced during cellular respiration, reacts with water:

CO+ H2O ↔ H2CO3 ↔ HCO3–+ H+

Notice that hydrogen ions are also generated, which increases the acidity of blood and decreases the pH. In your body, the hydrogen ions are absorbed by hemoglobin molecules on your red blood cells. Meanwhile, the bicarbonate ions circulate in the blood plasma, preventing rapid pH changes. As your blood circulates past the metabolizing cells, more and more CO2 enters your bloodstream and turns to bicarbonate ions. By the time the blood reaches the lungs, it is full of bicarbonate and hydrogen ions. The bicarbonate and hydrogen ions now combine and the reaction goes from right to left, releasing the CO2, which is now breathed out. In this exercise, you will bubble CO2 into tap water and demonstrate the change in pH as carbonic acid is formed. You will use the pH indicator phenol red, which turns yellow in acidic conditions and magenta (red-purple) in basic conditions.


  • Tap water
  • Drinking straw
  • Ehrlenmeyer flask (250 mL)
  • Phenol red
  • pH paper


  1. Obtain a small flask and a straw, and fill the flask approximately ¼ full with tap water.
  2. Measure the pH of the water using the pH paper.
  3. Add 6–7 drops of phenol red to the flask.
  4. Do not swirl the flask (this may introduce CO2 into the solution!), but agitate gently to mix the solution.
  5. Record the initial color of the water.
  6. Using a straw, blow air bubbles into the solution and observe any color changes.
  7. Record the final pH of the solution.


Initial Solution Final Solution

Lab Questions

  1. What happens when carbon dioxide combines with water?
  2. Why did the phenol red solution turn color after you blew air bubbles into it?
  3. If a person holds her breath, CO2 builds up in the bloodstream. What effect does this have on blood pH?
  4. If a person hyperventilates, too much CO2 is removed from the bloodstream. What effect does this have on blood pH?
  5. Why is “breathing into a bag” a good treatment for a hyperventilating patient?
  6. Why is pH homeostasis so critical in living organisms?

Part 4: Polar and Nonpolar Compounds

Screen Shot 2015-07-09 at 9.28.03 AMWater is a fascinating molecule whose chemical structure is pretty much responsible for life on earth. Chemical reactions that take place inside of a cell exist in an aqueous environment consisting principally of water. The primary characteristic of the water molecule that imparts its many unique qualities is the simple fact that water is a polar molecule. When considering the polarity of water, you must first remember that chemical bonds occur when 2 or more atoms “share” electrons. In the water molecule, the oxygen atom ‘hogs’ the electrons it shares with the hydrogen ions. Because the electrons are closer to the oxygen atom, that side of the water molecule ends up being partially negative while the hydrogen side of the molecule ends up being partially positive. This makes water an excellent solvent. Substances that are hydrophilic (love water) are usually polar or ionic molecules themselves, and will dissolve readily in water. Substances that are hydrophobic (hate water) are usually nonpolar molecules, and will not dissolve in water. Nonpolar substances will dissolve in a nonpolar solvent such as oil. Surfactants are special molecules that are both hydrophilic and hydrophobic. They allow water and oil to mix. Soaps and detergents are both examples of surfactants.


  • 2 test tubes
  • Oil
  • Tap water
  • Beet juice
  • Chili oil
  • Detergent


  1. Obtain two test tubes and add 5 ml of water and 5 ml of oil into each tube. Allow the tubes to stand for one minute. Record the appearance of the tubes and label the ingredients in the tube.
  2. Add ≈6 drops of beet juice extract to tube #1 and ≈6 drops of chili oil to tube #2. Allow diffusion to take place for 1–2 minutes. Record the appearance of the tubes.
  3. Shake each tube gently and let stand for several minutes. Record the appearance of the tubes.
  4. Next add a few drops of detergent to each tube; shake gently and observe. Record the appearance the tube.

Download this page to record the appearance of the tubes at every step.

Lab Questions

  1. What happens when lipids and water are combined? Why?
  2. How do beet juice extract and chili oil differ in their chemical properties? How do you know?
  3. Explain what happened when the tubes were shaken. What happened after the detergent was added? How can you explain these results?
  4. How is the phospholipid bilayer that makes up a cell membrane both hydrophilic and hydrophobic?
  5. What is a surfactant? How does it work?