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.
Boiling Point Determination
Similar to the melting point, the boiling point is a physical property. If the sample is a pure compound, then the boiling point can be used to determine the identity of the compound. Ultimately, experimentally determining the exact boiling point is challenging. Like melting points, experimental boiling points are given as a range and vary by a few degrees from the actual literature value.
Vapor Pressure
To understand why a solvent boils, which is characterized by the familiar bubbling of the solution, it is important to understand the dynamics between the liquid and gaseous phases. Consider a pure liquid compound in a sealed container. Some of the molecules at the surface of the liquid will have enough energy to overcome intermolecular forces and enter the gaseous phase. However, molecules in the gas phase may also lose energy and condense back into a liquid. Therefore, there are two competing processes in this system: evaporation and condensation.
When the rate of evaporation is equal to the rate of condensation, the system has reached a state of equilibrium. This means that for every molecule that enters the gas phase, another one condenses to the liquid phase, and there is no net gain or loss of the liquid or gas amount in the container. Once equilibrium has been established, the pressure exerted by the vapor above the liquid is called the vapor pressure. The tendency for a liquid to vaporize is called its volatility. A more volatile liquid has a higher vapor pressure, while a less volatile liquid has a lower vapor pressure.
Vapor pressure varies by temperature. If you increase the temperature of the solution, more molecules have enough energy to escape the liquid phase, and thus the vapor pressure increases. Ultimately, if enough heat is applied, molecules that are not at the interface between the liquid and gas will transition into the gas phase and form the familiar bubbles we associate with boiling.
The boiling point of a liquid is reached when the total vapor pressure of the liquid is equivalent to the atmospheric pressure. The temperature at which this occurs is called the boiling point. At higher elevations, and thus a lower atmospheric pressure, a liquid will boil at a lower temperature, as less heat is required to increase the vapor pressure to the atmospheric pressure. Additionally, the volatility, or the ability of a solvent to vaporize, also impacts the vapor pressure. Solvents with high volatility have higher vapor pressure than solvents with lower volatility.
Factors that Impact Boiling Point
The similarity between melting points and boiling points means that the same factors that impact the melting point of a compound will also impact the boiling point. Therefore, the strength and types of intermolecular forces that are found within the liquid compound will affect the boiling point. Recall that there are three types of molecular forces: hydrogen bonding, dipole-dipole interactions, and London Dispersion forces. Each of these has different attraction strengths and require different amounts of energy to overcome. Compounds that can hydrogen bond will have higher boiling points than compounds that can only interact through London dispersion forces. An additional consideration for boiling points involves the vapor pressure and volatility of the compound. Typically, the more volatile a compound is, the lower its boiling point.
The Capillary Method to Determine Boiling Point
A simple method for determining the boiling point of an organic compound is to utilize the capillary method. In this setup, an empty glass capillary tube is inverted into a container of the pure compound in the liquid phase. As the liquid is heated, the vapor pressure of the sample increases and gaseous vapor starts to enter the glass capillary tube. This forces out the air trapped inside and results in bubbles emerging from the bottom of the capillary tube. At this point, the liquid is allowed to cool. Once the vapor pressure of the sample is the same as the atmospheric pressure inside the glass capillary tube, liquid will begin to enter the tube. The temperature of the solution when this phenomenon occurs is the boiling point of the liquid compound.
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.
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.
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