What is the capillary method for determining boiling point?

The capillary method involves placing an inverted capillary tube in a liquid sample and heating it gradually. As temperature increases, air in the capillary escapes and is replaced by the liquid's vapor. When vapor pressure exceeds atmospheric pressure, bubbles stream from the capillary. The boiling point is recorded when the liquid begins refilling the capillary after cooling, indicating vapor pressure has dropped to atmospheric pressure.

Why does ethanol have a higher boiling point than acetone?

Ethanol has a higher boiling point due to hydrogen bonding between its molecules, which is a stronger intermolecular force than the dipole-dipole forces present in acetone. Hydrogen bonding allows ethanol molecules to hold onto each other more effectively, requiring more thermal energy to transition from liquid to gas phase. Although acetone has a slightly higher molecular weight, molecular structure and intermolecular forces have a greater impact on boiling point than mass alone.

How does molecular structure affect boiling point?

Boiling point is directly related to molecular structure through the strength of intermolecular forces. Stronger intramolecular forces, such as hydrogen bonding, result in higher boiling points because molecules remain bonded in the liquid phase longer. Molecular weight has less impact than structure; for example, ethanol has a lower molecular mass than butane but a much higher boiling point due to hydrogen bonding, while butane relies on weaker van der Waals forces.

What experimental errors can affect boiling point measurements?

Common experimental errors include heating the water bath too rapidly, which prevents accurate temperature equilibration, and poor alignment of the thermometer and sample. Rapid heating can cause temperature overshooting and inaccurate bubble observation. Misalignment prevents the thermometer from accurately reading the sample temperature. Careful temperature control and proper thermometer positioning are essential for obtaining boiling point values that match reported literature values.

How do you know when the boiling point has been reached during the experiment?

The boiling point is identified by two key observations. First, when heating, record the temperature at which a rapid and continuous stream of bubbles emerges from the capillary tube. Second, when cooling, observe when the liquid begins to rise back into the capillary as bubble production decreases. The temperature at which liquid refills the capillary represents the boiling point, as this marks when vapor pressure equals atmospheric pressure.

Why is thermometer alignment important in boiling point determination?

Proper thermometer alignment ensures the thermometer bulb is level with the liquid sample, allowing accurate temperature measurement of the boiling liquid. Misalignment causes the thermometer to read incorrect temperatures, leading to inaccurate boiling point values. Keeping the thermometer and sample close together and level prevents temperature gradients and ensures the thermometer responds to the actual temperature of the liquid being heated.

What happens to vapor pressure inside the capillary as temperature increases?

As temperature increases, vapor pressure inside the capillary rises proportionally. Initially, air in the capillary escapes and is replaced by the liquid's vapor. When vapor pressure exceeds atmospheric pressure, vapor escapes in a continuous stream of bubbles. Upon cooling, vapor pressure decreases until it falls below atmospheric pressure, at which point the liquid is drawn back into the capillary, marking the boiling point temperature.

Boiling Points

Concepts Instructor Prep Student Protocol

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Liquids are held together by weak interactions between their molecules. To transition to the gas phase, a molecule must gain enough energy from heat or other sources to overcome these interactions.

Let's consider a closed container of a pure liquid at room temperature. Some of the liquid phase molecules at the surface gain enough energy to evaporate, and some of that vapor loses energy and condenses back to liquid. When the rates of evaporation and condensation are the same, the liquid and vapor reach equilibrium. At this point, the space above the liquid is saturated with vapor, and there is no net change in the amount of liquid and vapor in the container. The pressure exerted by the vapor is called the vapor pressure.

Now, consider an open container of liquid. Here the vapor can escape, keeping the area above the liquid from being saturated. Thus, evaporation could continue until there is no liquid left. As the temperature of a liquid increases, so does the evaporation rate, and therefore, so does the vapor pressure. If enough heat is added, molecules deeper in the liquid start vaporizing, which we see as bubbles of vapor forming.

This is called boiling, and it starts when the vapor pressure of the compound is equal to the atmospheric pressure. The temperature at which this occurs is the boiling point. Since each pure substance has different types and strengths of intermolecular interactions, different substances have different boiling points.

Now, let's take a look at how intermolecular interactions impact the boiling point. Many intermolecular interactions involve areas with uneven electron density called dipoles. All molecules have brief fluctuations in their electron distribution, so even nonpolar molecules can temporarily have dipoles. In response, opposing dipoles form in nearby molecules.

The attractive forces between temporary and induced dipoles are called London dispersion forces, and they are the dominant interactions between nonpolar molecules like hydrocarbons. London dispersion forces are stronger between large molecules with many electrons because those molecules hold their electrons weakly, making it easier for neighboring molecules to affect their electron distribution.

For example, a high molecular weight unbranched alkane has a higher boiling point than a lower molecular weight unbranched alkane. Additionally, an unbranched alkane has a higher boiling point than a branched alkane with a comparable molecular weight due to its higher surface area.

Molecules with polar bonds have permanent dipoles, enabling regular dipole-dipole interactions between them. Dipole-dipole interactions are stronger than London dispersion forces, so it usually takes more thermal energy to boil polar liquids than nonpolar liquids. For example, the boiling point of one chloropentane, a polar molecule, is 108 °C. Hexane, which is similar in surface area but lacks a dipole, has a lower boiling point.

Molecules with both a hydrogen covalently bound to a more electronegative atom and an electron-withdrawing atom with a lone pair of electrons are capable of hydrogen bonding. Hydrogen bonds are stronger than dipole-dipole interactions and London dispersion forces, so it takes even more thermal energy to overcome them. For example, propane and ethanol have similar molecular weights and surface areas. But the boiling point of propane is much lower than ethanol's because ethanol molecules can form hydrogen bonds with each other, while propane cannot.

In this lab, you will use the capillary method to determine the boiling point of two organic solvents.

Print Procedure Steps

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Source: Lara Al Hariri and Ahmed Basabrain at the University of Massachusetts Amherst, MA, USA

Boiling point determination of acetone and ethanol
The temperature at which a pure organic substance changes from the liquid phase to the gas phase is known as the boiling point. A liquid's boiling point can be determined using the capillary method, where an inverted capillary is placed in the liquid of interest and the liquid is heated. As the temperature increases, the air in the capillary escapes and is replaced by the vapor of the liquid. The vapor pressure in the capillary increases with temperature. Once it exceeds the atmospheric pressure, the vapor escapes the capillary in a stream of bubbles. When the heat is removed, the liquid cools, and the vapor pressure in the capillary decreases. When the vapor pressure reaches the atmospheric pressure, the liquid begins to fill the capillary. The temperature at which this occurs is the boiling point.
- Put on a lab coat, splash-proof safety glasses, and nitrile gloves. This experiment must be conducted in a hood.
- Attach a small test tube to the thermometer using a rubber band, and secure the thermometer in the clamp.
- Obtain acetone and bring it back to your hood. Using your glass pipette and a bulb, measure 1 mL of acetone and transfer it into the small test tube.
- Align the thermometer end so that it is level with the acetone in the test tube, keeping them close to each other. Invert a capillary tube and place it in the test tube so that the open end is facing downwards.
- Use a 250-mL beaker to create a water bath. Add approximately 180 mL of water to the beaker, and place it on the hotplate.
- Lower the thermometer and test tube into the water bath. Turn on the hotplate to the lowest setting, about 30°C.
- Slowly increase the temperature on the hotplate by 10 – 20°C every 10 min, and closely observe the liquid in the test tube. When you start to see occasional bubbles in the liquid, increase the heat setting by only 5°C every 10 min.
- Closely observe the capillary tube inside the test tube. Report the temperature at which a rapid and continuous stream of bubbles comes out of the capillary.
  Table 1: Boiling point of acetone and ethanol
  
  Boiling point temperature (°C)
  Bubbles Liquid in capillary
  Acetone
  Ethanol
  Click Here to download Table 1
- Turn off the heat and allow the water bath to cool. Keep observing the capillary as the production of bubbles decreases until no bubbles emerge from the capillary. The liquid will start to rise in the capillary. Record the temperature at which this occurs.
- Label a clean 25-mL beaker as ‘organic waste’. Once the water bath has cooled to about 35°C, remove the thermometer and test tube from the bath. Detach the test tube from the thermometer and pour the acetone into the waste beaker.
- Use a clean test tube and capillary and repeat the experiment using 1 mL of ethanol.
- After you have measured the boiling point of ethanol, allow the water bath to cool. Once it has cooled sufficiently, remove the test tube and thermometer from the water bath, detach the test tube from the thermometer, and pour the ethanol into the organic waste beaker.
- Dispose of the water from the water bath down the sink and place the capillaries and test tubes in the glass disposal. Dispose of the organic waste in the container provided by your instructor.
- Wash all of your glassware with detergent and water.
Results
In this experiment, we measured the boiling point of acetone to be 56°C, which compares well to the reported value. Similarly, the boiling point of ethanol was measured to be 78°C. Errors in the boiling point measurement can be attributed to many experimental errors, such as heating the water bath too rapidly, or poor alignment of the thermometer and sample.

The boiling point of an organic substance is directly related to its structure, where stronger intramolecular forces result in a higher boiling point as molecules are able to hold onto each other and remain in the liquid phase longer. The higher boiling point for ethanol is observed due to the OH structure that causes hydrogen bonding between the molecules. Acetone has a polar CO double bond, which results in dipole-dipole forces. Since hydrogen bonding is stronger than dipole-dipole forces, ethanol has a higher boiling point.

Additionally, ethanol has a lower molecular weight than acetone. However, molecular weight has less of an impact on the boiling point than the molecular structure. For example, butane is a gas at room temperature and pressure, as it has a boiling point lower than 25°C. Ethanol has a slightly lower molecular mass than butane, but it is liquid at room temperature and, therefore, has a boiling point higher than room temperature. This is due to the hydrogen bonding between the ethanol molecules, which is stronger than the van der Waals forces between the butane molecules.