10.3 Phase Transitions

Many medical tests require a small amount of blood to be drawn, for example to determine the hematocrit level in an athlete.

capillary action, the ability of a liquid to flow up a small tube against gravity, can be used to perform this procedure.
When the open end of a narrow-diameter glass tube touches the drop of blood, the forces between the molecule in the blood and those at the glass surface draw the blood up the tube.
The diameter of the tube and the type of fluid affect how far up the tube the blood goes.
A small tube has a relatively large surface area for a given volume of blood, which results in larger attractive forces, allowing the blood to be drawn farther up the tube.
The liquid is held together by its own forces.
The liquid stops rising when the downward force equal to the upward force associated with capillary action is equal to the weight of the tube.

capillary action draws blood into a small glass tube for analysis.

Changes of physical state, or phase transitions, are witnessed and utilized in a number of ways.

One example of global significance is the melting of water.
These changes of state are important parts of our earth's water cycle as well as many other natural phenomena and technological processes of central importance to our lives.
The essential aspects of phase transitions are explored in this module.

Gas can't escape when a liquid is in a closed container.

The gas phase molecule will sometimes collide with the surface of the Condensed phase, and in some cases it will result in the molecule reentering the Condensed phase.
This is not a static situation, as the molecules are constantly exchanged between the two phases.
The size of the vessel and the surface of the liquid in contact with the Vapor have no effect on the pressure of the Vapor.
We can measure the vapor pressure of a liquid by placing a sample in a closed container and using a manometer to measure the increase in pressure that is due to the vapor in equilibrium with the Condensed phase.

Dynamic equilibrium is reached when the rate of gas escaping from the liquid increases and the rate of gas entering the liquid equals the rate of gas entering the liquid.

The vapor pressure of the gas is constant when this equilibrium is reached.

Different substances will exhibit different equilibrium vapor pressures because the chemical identities of the molecule in a liquid determine the types and strengths of intermolecular attractions.

Strong intermolecular attractive forces will serve to impede vaporization as well as favoring "recapture" of gas-phase molecules when they collide with the liquid surface, resulting in a relatively low vapor pressure.
Weak intermolecular attractions present less of a barrier to vaporization and a reduced likelihood of gas recapture.
The dependence of vapor pressure on attractive forces is shown in the following example.

Most of the attractions of diethyl ether are London forces.

Although this molecule is the largest of the four under consideration, its IMFs are the weakest and, as a result, its molecule most readily escape from the liquid.
It has the highest pressure.
Diethyl ether has stronger dispersion forces than ethanol.
Diethyl ether has a higher vapor pressure than ethanol because it has a lower ability to hydrogen bond and because fewer molecules escape from the liquid at any given temperature.

Liquids and Solids hydrogen bonding provides stronger intermolecular attractions, fewer molecules escaping the liquid, and a lower vapor pressure than for either diethyl ether or ethanol.

Like water, ethylene glycol has two -OH groups.
It is larger than water and experiences larger London forces.
Its overall IMFs are the largest of the four substances, which means its vaporization rate will be slower and its pressure will be the lowest.

These compounds exhibit hydrogen bonding and are difficult to overcome, so the vapor pressures are relatively low.

The vapor pressures decrease as the size of molecule increases from methanol to butanol.

The average ke of a liquid increases as temperature increases.

At any given temperature, the molecule of a substance has a range of energy with a certain percentage of it having enough to escape the liquid.

The higher the average speed of the molecule that escapes, the higher the vapor pressure.

The distribution of energy for the molecule in the liquid is affected by temperature.

More molecule have the necessary energy to escape from the liquid into the gas phase.

The liquid reaches its boiling point when the atmospheric pressure is equal to the vapor pressure.

The pressure is due to the atmosphere.
The dependence of a liquid's boiling point on surrounding pressure may be depicted in the curves.

The boiling points of liquids are the temperatures at which their equilibrium vapor pressures equal the pressure of the surrounding atmosphere.

The elevation of Leadville, Colorado is 10,200 feet.

The atmospheric pressure in Leadville will be equal to the vapor pressure in water.

The boiling point of ethyl ether was measured at a base camp on the slopes of Mount Everest.

Figure 10.24 shows the atmospheric pressure at the camp.

This linear equation can be expressed in a two-point format that is convenient for use in various computations, as shown in the example exercises that follow.

The octane rating of 2,2,4-trimethylpentane is 100.

It is one of the standards used for the octane-rating system.
The vapor pressure of isooctane is 100.0 kPa at 34.0 degC and 98.6 degC.
This information can be used to estimate the enthalpy of vaporization.

At 20.0 degC, the vapor pressure of the substance is 52.95 kPa, and at 63.5 degC, it is 53.3 kPa.

This information can be used to estimate the enthalpy of vaporization.

The normal boiling point for benzene is 80.1 degC and the enthalpy of vaporization is 30.8 kJ/mol.

The normal boiling point for acetone is 56.5 degC and the enthalpy of vaporization is 31.3 kJ/mol.

Vaporization is an endothermic process.

The cooling effect can be seen when you leave a shower.
You feel cold when the water on your skin evaporates.

The reverse of an endothermic process is called exothermic.

evaporation of the water in sweat is one way our body is cooled.

We can lose 1.5 L of sweat per day in very hot climates.
Assuming that sweat is pure water, we can get an approximate value of the amount of heat removed by evaporation.

Evaporation of sweat cools the body.

We start with the volume of sweat and use the given information to convert to the amount of heat needed: 1.5 L x 1000 g 1L x 1 mol 18 g x 43.46 kJ

When we heat a solid, we increase its average energy, which makes it hotter.

At this point, the temperature of the solid stops rising, despite the constant input of heat, and it remains constant until all of the solid is melted.

A small amount has melted.

The ice does not change its temperature.

The solid and liquid phases remain in equilibrium if we stop heating during melting and place the mixture in a perfectly insulated container.

It's like this with a mixture of ice and water in a thermos bottle, with almost no heat getting in or out, and the mixture of solid ice and liquid water remains for hours.
The direction of the phase transition being considered affects the use of one term or the other.

The strength of the attractive forces between the units in the crystal affects the enthalpy of fusion and the melting point.

The weak attractive forces of the Molecules form low melting points.

At higher temperatures, particles with stronger attractive forces melt.

The enthalpy of fusion is the amount of heat required to change one mole of a substance from the solid state to the liquid state.

6.0 kJ/mol is the enthalpy of fusion of ice.

At room temperature and standard pressure, a piece of dry ice appears to disappear without ever forming a liquid.

At the melting point of water, snow and ice form a slow process that may be accelerated by winds and the reduced atmospheric pressures at high altitudes.
There is a vivid purple vapor when solid iodine is warmed.
The formation of frost is an example of deposition.

A purple gas is produced in the bottom of the tube and deposited on the colder part of the tube above.

Vaporization requires an input of energy to overcome intermolecular attractions.

The energy required to convert one mole of a substance from a solid to a gaseous state is called the enthalpy.

The attractions need to be partially overcome in order to convert a solid into a liquid.

The enthalpy of fusion is less than the enthalpy of vaporization.
This same logic can be used to derive an approximate relation between the phases of a substance.
Sublimation may be modeled as a two-step process of melting followed by vaporization in order to apply Hess's Law.

The enthalpy of fusion and enthalpy of vaporization is the same for a given substance.

The relation does not apply to matter being heated or cooled.

When a substance reaches a temperature corresponding to one of its phase transitions, further gain or loss of heat is a result of diminishing or enhancing intermolecular attractions.
The substance's temperature remains constant as it undergoes a change in state.

A pot of water can be boiled.

The heat from a stove burner will increase the water's temperature initially.
When the water reaches its boiling point, the temperature remains constant despite the continued input of heat from the stove burner.
The same temperature is maintained if the water is boiling.
The water temperature does not rise if the burner setting is increased to provide more heat.
For other phase transitions, this behavior is also observed: For example, temperature remains constant progress while the change of state is in.

A heating curve is a depiction of temperature changes in a substance as it absorbs more heat.

When a substance undergoes phase transitions, there are regions of constant temperature in the curve.