Wet Dry Ice Lab

Phase Changes and Phase Diagrams

Introduction

As dry ice sublimes in a closed system, its three phases are clearly viewed and its phase diagram takes on concrete meaning

to students.

Concepts

• Phase changes • Triple Point

• Phase diagrams • Sublimation

Materials

Safety Precautions

Dry ice is extremely cold, handle only with insulated heavy cloth gloves and never with wet hands; frostbite is possible with only brief

exposure. The demonstrator and all observers must wear chemical splash goggles. Do not attempt this demonstration on a larger scale in

plastic bottles or other containers—the plactic may shatter violently and cause injury. Please review current Material Safety Data Sheets

for additional safety, handling, and disposal information.

Procedure

1. Pulverize the dry ice into small pieces about the size of rice grains or sugar crystals. Observe that the dry ice does not melt,

it sublimes. The resulting “fog” is due to water vapor condensing on the extremely cold CO

gas that is produced.

2. Cut off the tapered end of a wide-stem, Beral-type pipet. Scoop about 8–10 pieces of dry ice into the stem of the pipet

the same time (solid, liquid, and gas).

6. Release the grip on the pliers to relieve some of the pressure in the pipet. A loud pop is produced and the CO

immediately

returns to a solid—the dry ice looks like fluffy snow.

7. Repeat the demonstration by folding the open end of the pipet stem over again and reclamping the pipet shut with the

pliers. The CO

repeated 3–4 times.

8. Your students will undoubtedly ask: “What will happen if you don’t release the pressure?” Repeat the demonstration, but

don’t release the pressure when the liquid (“wet dry ice”) begins to boil. Caution: Everyone should be wearing goggles as

Disposal

Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review

all federal, state and local regulations that may apply, before proceeding. Allow the dry ice to sublime in a well-ventilated area,

such as a hood, and the used pipets may be placed in the trash according to Flinn Suggested Disposal Method #26a.

Tips

• This activity is written as a demonstration, but it is safe enough to be used as a student activity. Monitor students so they

do not add too much dry ice to their pipets and strictly enforce the “wear your goggles rule.”

Dry ice is usually a fascinating and fun material for your students. From making “fog” to “boiling in water,” it is well-known for

creating special effects. Carbon dioxide, however, also has fascinating and very useful chemical properties. At room temperature

and pressure, solid carbon dioxide will warm to –78 °C and then begin to sublime to carbon dioxide gas. The carbon dioxide gas

is, initially, also at –78 °C, which causes moisture in the air to condense and form the characteristic fog that dry ice is famous for.

One interesting feature of carbon dioxide is that at atmospheric pressure, it only exists as a solid or gas. In order to exist as a

liquid, carbon dioxide must be subjected to a pressure of at least 5.11 atmospheres. 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 below).

A phase diagram has temperature as the independent (x) and pressure as the dependent (y) axis. Three distinct regions are rep-

resented as regions of pressure and temperature relative to the state of the substance as solid, liquid, or gas. The boundaries

pressure (1 atm. or 14.7 psi) and room temperature (25 °C), the solid will sublime spontaneously, maintaining its temperature of

–78 °C (point ①) until the solid disappears completely. The gas formed will absorb heat from the room until it obtains a stable

gaseous state at 25 °C. If a sample of dry ice is sealed in a closed system such as the triple point apparatus, the pressure begins

to rise allowing the solid to exist at a higher temperature. Most of the energy absorbed from its surroundings will result in an

increased temperature and corresponding increase in vapor pressure. The effect of this increased energy is a series of pressure–

temperature equilibria along the solid–gas line between points ① and ② on the phase diagram.

Once the triple point is reached (point ②), the energy absorbed causes the solid to melt. The solid–gas and liquid–gas phases are

also in equilibrium. While this phase change occurs, the temperature and pressure remain constant at 5.1 atm. and –56.6 °C as

long as the solid, liquid, and gaseous phases are in contact with each other. If the solid becomes covered with liquid, an

equilibrium no longer exists between all three phases, and both the temperature and pressure will increase into the liquid

CO

2

(l) → CO

2

(g) ∆H = 16 kJ/mol

CO

2

(l) → CO

2

(s) ∆H = –9 kJ/mol

This demonstration also presents a good opportunity to discuss the various units that are used when measuring pressure.

Connecting to the National Standards

This laboratory activity relates to the following National Science Education Standards (1996):

Unifying Concepts and Processes: Grades K–12

Systems, order, and organization

Evolution and equilibrium

The pressure inside the pipet rose and the carbon dioxide began to sublime, turning into a gas.

b. When the pressure stabilized

There is carbon dioxide present in all three states, solid, liquid, and gaseous.

c. When the solid CO

2

The pressure began to rise again.

d. When the pressure in the pipet was released

The liquid carbon dioxide froze back into its solid state.

2. Explain how the triple point was reached within the pipet.

When the dry ice was sealed in the pipet, the pressure within the pipet began to rise. Solid carbon dioxide, which in normal room