Boiling in a Syringe

Introduction

Water boils at 100 °C, right? Not always! Can water boil at room temperature? Explore a phase change diagram to determine why there are different directions for food preparation in higher altitudes, what a pressure cooker does and how water may boil in a syringe even if it is not 100 °C.

Concepts

• Vapor pressure
• Boiling point
• Phase change
• Phase diagram

Materials

Safety Precautions

Be careful when pulling the plunger back. The plunger may snap back very quickly if it is not held tightly. Use caution when using a hot plate or handling hot materials. Wash hands thoroughly with soap and water before leaving the laboratory. Follow all laboratory safety guidelines.

Disposal

The water may be disposed of down the drain. Save the syringe and stopcock for future use.

Prelab Preparation

1. In a borosilicate beaker, heat tap water to 80–90 °C. Do not allow the water to boil.
2. Fill a second beaker with warm tap water ranging between 40 and 50 °C.
3. Insert the stopcock into the end of the syringe and twist counterclockwise until tight. 

Procedure

1. Ensure the stopcock is open (see Figure 1) and draw up approximately 25 mL of hot water into the syringe.
   {12294_Procedure_Figure_1_Open}
2. Hold the syringe so the stopcock is pointed up and carefully adjust the volume of the syringe to remove any air bubbles.
3. Close the stopcock (see Figure 2).
   {12294_Procedure_Figure_2_Closed}
4. Pull back on the plunger and observe the water.
5. Eject the water from the syringe and repeat the activity with warm tap water.

Student Worksheet PDF

12294_Student1.pdf

Teacher Tips

• Experiment with different water temperatures or even different liquids—such as ethyl alcohol. Note: Please refer to current Safety Data Sheets for relevant safety, handling and disposal information.
• The lower the temperature, the lower the “external” or applied air pressure inside the syringe needs to be to make the water boil. This means the syringe plunger would have to be pulled further out from the liquid, which reduces the overall pressure inside the syringe.
• The water in the syringe will boil rapidly at first and then the boiling will decrease as the pressure inside the syringe increases from the water that has already boiled and entered the vapor phase. If desired, release the pressure on the syringe momentarily and then resume the demonstration.
• A video of this demonstration, Boiling in a Syringe, presented by DeWayne Lieneman, is available for viewing as part of the Flinn Scientific “Teaching Chemistry” eLearning Video Series. Please visit the eLearning website at http://elearning.flinnsci.com for viewing information. This video is part of the Evaporation and Boiling video package.

Answers to Questions

1. Record all observations.
   {12294_Answers_Table_1}
2. What happened after the syringe plunger was pulled back? Explain.

   The water began bubbling rapidly because the pressure decreased which also decreased the boiling point of the liquid.

3. Why does the observation described above slow down after 45 seconds?

   The vapor pressure increases from the water that has already boiled and entered the vapor phase.

4. The following graph represents the phase diagram of water.
   {12294_Answers_Figure_3}
   1. See graph.
   2. See graph.
   3. See graph.
   4. See graph.
   5. Looking at data point B—if the pressure remains constant, what would need to happen for the water to return to a completely liquid state? Mark a new data point C representing this action.

      To completely return to a liquid state at constant pressure, the temperature would have to be reduced.

5. Give an example how pressure affects the boiling point of water and provide a practical example.

   Student answers will vary. May include high altitude cooking or pressure cooking.

Discussion

Every liquid boils at the temperature at which its vapor pressure equals the pressure above its surface. By decreasing the pressure inside the syringe, water will boil below 100 °C. When the plunger is originally pulled, the air pressure in the syringe falls below the water’s vapor pressure, causing the water to boil. If the plunger is held back long enough, the boiling slows and eventually stops. As the water boils, the water vapor produced is pressurizing the area above it. The water will continue to boil until the pressure equals that of the vapor pressure or until there is no liquid left—whichever comes first.

Water boils at 100 °C at 1 atmosphere of pressure—this is called the normal boiling point. Phase changes are temperature and pressure dependant. Every liquid boils at the temperature at which its vapor pressure, the force per area exerted by the vapor, equals the pressure above its surface. Change the pressure and the temperature needed to boil changes too! At a higher altitude the atmospheric pressure is lower than at sea level. See the boiling point of water at various locations in the table.

At a constant pressure a phase change diagram for water is a traditional heating or cooling curve as shown in Figure 4. Note: Heating and cooling curves are constructed under the assumption that pressure does not change and are traditionally graphed with temperature on the y-axis and either time under constant heating or heat energy on the x-axis. At point A water is in its solid form—ice. As heat is added the temperature of the ice increases by increasing kinetic energy of the the H₂O(s) molecules until the ice begins to melt—point B. From points B to C heat is still being added to the H₂O. In this range, the temperature does not increase but instead the heat energy is being used to convert the H₂O(s) to H₂O(l). Once all the ice has been converted to water (point C), the temperature of the water will begin to rise again. This temperature will continue to rise linearly until point D is reached where H₂O(l) will begin using the heat energy to convert to H₂O(g). Once again the temperature will plateau as the heat energy is used to convert from the liquid phase to the vapor phase—points D to E. Note: The phase change from liquid to steam takes about seven times as much energy as the conversion of ice to liquid water. At point E, when all of the liquid water has been converted to steam, the temperature of the steam will begin to rise. The heat energy being added to the system is no longer being used to convert phases and again is used to increase the kinetic energy and thus the temperature of the H₂O(g) molecules.

Most chemicals will exist as a solid, liquid or gas depending on temperature and pressure. This relationship between phase, pressure, and temperature can be presented graphically in the form of a phase diagram (see Figure 5). A phase diagram has temperature as the independent (x) and pressure as the dependent (y) axis. Three distinct regions are represented as regions of pressure and temperature relative to the state of the substance as solid, liquid, or gas. The boundaries between regions show the values of pressure and temperature when two phases are in equilibrium. For example, sublimation/deposition occurs at the boundary between solid and gas, vaporization/condensation occurs at the liquid–gas boundary, and melting/freezing occurs at the solid–liquid boundary. The point at which all three phase boundaries meet is called the triple point and signifies the temperature and pressure at which all three phases exist and are in equilibrium. How might pressure changes be observed outside of a lab setting? Pressure changes affect everyday activities such as cooking—both at higher pressures and lower pressures. A pressure cooker is a pot with a lid that seals with gaskets or by mechanical means. As the pressure cooker is heated, the pressure inside the cooker increases to more than atmospheric pressure. Remember, as the pressure decreased in the Boiling-in-a-Syringe demonstration, the boiling point also decreased. The opposite is true in a pressure cooker. As the pressure increases, the boiling point of water increases. Most pressure cookers increase the pressure by 15 psi which increases the boiling point of water from 100 to 122 °C. The higher temperature causes the food to cook faster—⅓ of the normal cooking time. At higher altitudes, the atmospheric pressure is lower so foods take longer to cook. Many times liquids also evaporate faster, so besides varying cooking times, recipe adjustments might be needed. Gases also expand more due to the lower pressure, so baking dough would rise faster in higher altitudes.