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Phase Change Notes
I. Phases of matter and phase changes
a. Recall the three main phases of matter:
Energy is involved during the transition from one phase of matter to another. You should learn how energy is involved in each of these phase transitions, and you should be able to name each of these phase transitions.
(1) Melting: energy absorbed
(2) Freezing: energy released
(3) Condensation: energy released
(4) Vaporization: energy absorbed
(5) Deposition: energy released
(6) Sublimation: energy absorbed
The heat content of a chemical system is called the enthalpy (symbol: H)
The enthalpy change (H) is the amount of heat released or absorbed when a chemical reaction occurs at constant pressure.
Figure 1: A pictorial summary of phase changes. For our purposes, "enthalpy" (symbol = H) is added heat energy.
II. Heat and temperature
a. Heat is a form of energy that flows from hot objects to cold objects. Heat is measured in joules.
b. Temperature is a measure of the average kinetic energy of the particles in a substance. Temperature is measured in °C or K.
c. Although it seems counterintuitive, a substance that is heated will not undergo a continuous rise in temperature. Likewise, if a substance is cooled, its temperature will not decrease uniformly.
In fact, as a substance such as ice is heated, its temperature will only increase until a phase change occurs. During a phase change, the temperature of a substance does not change.
When the temperature of a substance that is being cooled or heated does not change, it is undergoing a phase change.
d. In the diagram Heating/Cooling Curve for H20, one can see that as ice is heated from -10°C, its temperature increases only until 0°C. From -10°C to 0°C, all of the heat absorbed by the ice goes into increasing the kinetic energy of the ice molecules (they vibrate faster).
e. At 0°C, however, the particles of ice can not move any faster. Therefore, the temperature does not rise. At this point, any absorbed energy goes into making bonds between water molecules break. These bonds between water molecules are intermolecular. This period is called the Heat of Fusion.
i. The amount of energy needed to melt a substance at its melting point is called the heat of fusion of that substance. For H2O, that amount of energy is 6.02 kJ/mol.
ii. The temperature at which melting happens is called the normal melting point of the substance. Believe it or not the melting point is the same as the freezing point. They are the same temperature.
iii. The amount of energy that must be released in order for a substance to freeze at it s freezing point is also called the heat of fusion of that substance. For H2O, that amount of energy is 6.02 kJ/mol.
f. After the ice has entirely melted, the water that forms can be heated to a higher temperature. The temperature of water (at standard pressure) can only reach 100°C.
i. The temperature at which a substance boils is called the boiling point of as substance.
ii. The temperature at which a substance boils when the pressure is 1 atm (standard pressure) is called the normal boiling point of as substance. The boiling point and melting point of a substance depend on vapor pressure as well as temperature.
iii. The amount of heat needed to boil a substance at its boiling point is called the heat of vaporization of that substance. For H2O, that amount of energy is 40.6 kJ/mol.
iv. The amount of heat that must be released by a gas - at its boiling point – so that it can condense to a liquid is also called the heat of vaporization of that substance. For H2O, that amount of energy is 40.6 kJ/mol. The boiling temperature for a liquid must be the same as the condensation temperature for the gas phase of that substance.
III. Calculations involving the heating/cooling curve.
a. There are five different regions to the heating/cooling curve for water.
i. The heating/cooling curve for other substances are constructed the same for other substances, but will have different slopes and runs
ii. Water is a very important substance, it is the essence of life that without which the world as we know would not be. Therefore, knowing how to perform the calculations for water’s phase changes is very important to a scientist.
b. The heating/cooling of ice: Use q=mCΔT
c. The melting/freezing of H2O: Use 6.02 kJ/mol
d. The heating./cooling of water: Use q=mCΔT
e. The boiling/condensation of H2O: Use 40.6 kJ/mol
f. The heating/cooling of steam: Use q=mCΔT
IV. Intermolecular forces of attraction
a. These are the forces that hold molecule to molecule, not atom to atom.
b. Molecule to molecule = intermolecular; atom to atom in a molecule = intramolecular. Remember the three kinds of intramolecular bonds: ionic bonds ( ), covalent( ) or metallic bonds ( ).
c. We care about three main intermolecular forces of attraction: hydrogen bonding, dipole-dipole interactions, and London dispersion forces.
i. VERY STRONG: Hydrogen bonding – the attraction between an unshared pair on an atom of one molecule for an electropositive (“electronstarved”) hydrogen atom on another nearby molecule.
ii. SOMEWHAT STRONG: Dipole-dipole interactions – the attraction between to polar molecules
iii. WEAKEST: London dispersion forces aka Van Der Waals– the attraction that one nonpolar molecule has for another nonpolar molecule, due to the temporary uneven distribution of electrons within molecules which causes them to be weakly and temporarily polar.
V. Vapor pressure, equilibrium, and evaporation
a. All liquids evaporate when left out “in the air.” The highest energy molecules in the liquid phase, when they are on the surface of the liquid, escape to the gas phase.
b. Even contained liquids will evaporate, that is, they will proceed from the liquid to the gaseous state of matter.
c. However, in a contained liquid, as the liquid particles evaporate, the gas formed by them (the vapor) begins to exert a pressure just like any other gas in a container would.
d. As the pressure increases due to the growing number of gas particles, the likelihood of two gas particles interacting and re-liquefying (condensation) increases. After a while, the number of particles entering the vapor phase equals the number of particles re-entering the liquid phase. At this point, we say that a state of equilibrium has been established between the gas and liquid phases. Equilibrium is when maximum concentrations possible for a substance in all phases is achieved.
e. The pressure exerted by a gas that is in equilibrium with its own liquid phase at a certain temperature is called the vapor pressure of that substance.
f. Huge tables are published in huge books detailing the vapor pressures of substances. Alternatively, there are mathematical equations that can predict (or at least estimate) the vapor pressure of a liquid from some of its physical properties.
VI. Phase Diagrams
a. A particular substance can exist in several different phases – sometimes simultaneously – depending on the pressure and temperature at which the substance is held.
b. The phase state behavior of a substance can be summarized with a phase diagram.
c. Key features:
i. Triple point: the conditions of temperature and pressure at which all three phases of matter can coexist.
ii. Vapor pressure: the pressure at which a contained liquid is in equilibrium with its gas (vapor) phase. Use the G-L line two different ways:
1. The boundary line between the “gas” and “liquid” sections of the graph represents the vapor pressure of the substance. As temperature goes up, the vapor pressure goes up. So, you can determine the vapor pressure for a substance at a given temperature.
2. The boundary between G and L also represents the boiling temperature of the substance at a given pressure.
iii. Normal boiling point: this is the boiling point at standard atmospheric pressure (1 atm). Notice that water can be boiled at many temperatures, depending on the pressure. In Denver, Colorado, water boils at 95°C; on Mount Everest, water boils at 70°C. What non-scientists call “boiling point” is what we call “normal boiling point”.
iv. Normal freezing point: This is the freezing point at standard atmospheric pressure (1 atm). Notice that water can be turned to ice at many temperatures, depending on the pressure
d. Different substances have different phase diagrams because they behave differently as the pressure and temperature change.
Test your knowledge of phase diagrams by answering the following questions:
What is the triple point of H2O? What is the triple point of CO2?
What is the boiling temperature of water at 1 atm (101.3 kPa) of pressure?
Given the phase diagram for water state what phase(s) of water is/are present at each of the following temperature-pressure conditions:
a) at any point on curve BO
b) at any point on curve OA
c) at any point on curve OC
d) at point O
2. Based on the phase diagram from the previous question, what effect would each of the following changes have on a sample of water at any point on curve OA:
a) increasing the temperature at constant pressure
b) decreasing the pressure at constant temperature
c) decreasing the temperature at constant pressure
d) increasing the pressure at constant temperature
3. Given a sample of water at any point on curve BC in the phase diagram given above, how could more liquid water in the sample be converted into a solid at room temperature?
Practice Questions
I. Label each of the 5 phase changes in the diagram below with the letter of the correct response. Not all of the letters get used.
a. Vaporization
b. Sublimation
c. Melting
d. Freezing
e. Condensation
f. Deposition
II. Label each of the blanks below as either “energy required” or “energy released”.
A= ENERGY REQUIRED B= ENERGY RELEASED
III. The diagram below shows three different flasks. The substance in the flasks is H2O. Match each description with the letter of the appropriate flask in the diagram.
10) Which of the flasks shows a liquid that has reached equilibrium with its vapor?
11) Which of the flasks depicts a liquid that will eventually reach equilibrium with its vapor?
12) Which of the pictures shows a flask in which the liquid WILL NOT reach equilibrium with its vapor?
IV. The atmospheric pressure in Denver, Colorado on a particular day is 658 mm Hg. The table below lists the vapor pressure of water at various temperatures. Given that at 760 mmHg water boils at 100 °C determine the approximate temperature at which water will boil in Denver, CO on this day.
13) a) 100 °C b) 105 °C c) 96 °C d) 106 °C e) 0 °C
14. Order the intermolecular forces (dipole-dipole, London dispersion, and hydrogen bonding) from weakest to strongest.
[A] dipole-dipole, London dispersion, hydrogen bonding
[B] London dispersion, dipole-dipole, hydrogen bonding
[C] hydrogen bonding, dipole-dipole, London dispersion
[D] London dispersion, hydrogen bonding, dipole-dipole
15. The intermolecular forces called hydrogen bonding will not exist between molecules of
[A] NH3 [B] H2 [C] HF [D] H2O [E] any of these
16. At 1 atm of pressure and a temperature of 0°C, which phase(s) of H2O can exist?
[A] ice and water vapor [B] water only [C] ice only [D] ice and water [E] water vapor only
17. The normal freezing point of water is
[A] 0°F [B] 32°C [C] 273 K [D] 373°C [E] none of these
18. The normal boiling point of water is
[A] 373 K [B] 0°F [C] 32°F [D] 273 K [E] none of these
19. Calculate the quantity of energy required to change 3.00 mol of liquid water to steam at 100°C. The molar heat of vaporization of water is 40.6 kJ/mol.
[A] 300 kJ [B] 13.5 kJ [C] 122 kJ [D] 40.6 kJ [E] none of these
20. Calculate the quantity of energy required to change 26.5 g of liquid water to steam at 100°C. The molar heat of vaporization of water is 40.6 kJ/mol.
[A] 1.08 × 103 kJ [B] 59.8 kJ [C] 1.53 kJ [D] 27.6 kJ [E] none of these
21. The specific heat capacity of liquid water is 4.18 J/g°C. Calculate the quantity of energy required to heat 10.0 g of water from 26.5°C to 83.7°C.
[A] 572 J [B] 837 J [C] 239 J [D] 2.39 × 103 J [E] none of these
22. The molar heat of fusion of water is 6.02 kJ/mol. Calculate the energy required to melt 46.8g of water.
[A] 6.02 kJ [B] 7.77 kJ [C] 282 kJ [D] 2.32 kJ [E] none of these
23. The freezing point of helium is approximately −270°C. The freezing point of xenon is −112°C. Both of these are in the noble gas family. Which of the following statements is supported by these data?
a) The London forces between the helium molecules are greater than the London forces between the xenon molecules.
b) The London forces between the helium molecules are less than the London forces between the xenon molecules.
[C] As the molar mass of the noble gas increases, the freezing point decreases.
[D] Helium and xenon form highly polar molecules.
[E] none of these
24. Choose the state of water in which the water molecules are farthest apart on average.
[A] ice (solid) [B] steam (vapor) [C] all the same [D] liquid
25. The process of evaporation happens when which of the following occurs?
[A] A solid becomes a gas.
[B] A liquid becomes a solid.
[C] A solid becomes a liquid.
[D] A gas becomes a liquid.
[E] A liquid becomes a gas.
26. Which of the following processes must exist in equilibrium with the evaporation process when a measurement of vapor pressure is made?
[A] fusion [B] condensation [C] vaporization [D] boiling [E] sublimation
27. The vapor pressure for water at 100.0°C is
[A] 760 torr [B] More information is needed. [C] 85 torr [D] 1 torr [E] 175 torr