Chemistry: Phases and Phase Changes

Question 1

Gas

Answer 1

Conforms to the volume and shape of the container it is in. Continual motion. Low density. Easily compressed to smaller volume.

Question 2

Liquid

Answer 2

Conforms to the shape of the container; however, has definite volume. Sliding motion of particles past one another. Moderate density. Small ability to be compressed.

Question 3

Solid

Answer 3

Definite volume and shape. Particles in a fixed position. High density. Difficult to compress.

Question 4

Solutions

Answer 4

One of the most important properties of liquids is their ability to mix, both with each other and with other phases, to form solutions.

Question 5

Miscibility

Answer 5

Degree to which two liquids can mix.

Question 6

Immiscible

Answer 6

When two molecules tend to repel each other due to their polarity difference. Like oil and water.

Question 7

Emulsions

Answer 7

Under extreme conditions, such as violent shaking, two immiscible liquids can form a fairly homogenous mixture called an emulsion. Though they look like solutions, emulsions are actually mixtures of discrete particles to small to be seen distinctly.

Question 8

Crystalline Solid

Answer 8

Possesses an ordered structure; its atoms exist in a specific, 3D geometric arrangement with repeating patterns of atoms, ions, or molecules. Example is NaCl. Two most common forms of crystals are metallic and ionic crystals.

Question 9

Amorphous Solid

Answer 9

Has no ordered 3D arrangement, although the molecules are also fixed in place. Example is glass.

Question 10

Ionic Solids

Answer 10

Aggregates of positively and negatively charged ions; there are no discrete molecules. The physical properties of ionic solids include high melting points, high boiling points, and poor electrical conductivity in the solid phase. These properties are due to the compounds’ strong electrostatic interactions, which also cause the ions to be relatively immobile. Ionic structures are given by empirical formulas that describe the ratio of atoms in the lowest possible whole numbers.

Question 11

Metallic Solids

Answer 11

Consist of metal atoms packed together as closely as possible. Metallic solids have high melting and boiling points as a result of their strong covalent attractions. Pure metallic structures (consisting of a single element) are usually described as layers of spheres and roughly similar radii.

Question 12

Unit Cells

Answer 12

The repeating units of crystals (both ionic and metallic) are represented by unit cells. There are many types of unit cells. Atoms are represented as points, but are actually adjoining spheres. Each unit cell is surrounded by similar units. In the ionic unit cell, the spaces between points (anions) are filled with other ions (cations).

Question 13

Evaporation

Answer 13

The temperature of a liquid is related to the average kinetic energy of the liquid molecules; however, the kinetic energy of the molecules will vary. A few molecules near the surface of the liquid may have enough energy to leave the liquid phase and escape into the gaseous phase (evaporation/vaporaization). Each time the liquid loses a high-energy particle, the temperature of remaining liquid decreases; thus, evaporation is a cooling process. Given enough kinetic energy, the liquid will completely evaporate.

Question 14

Condensation

Answer 14

If a cover is placed on a beaker of liquid, the escaping molecules are trapped above the solution. These molecules exert a countering pressure, which forces some of the gas back into the liquid phase; this process is called condensation. Atomospheric pressure acts on a liquid in a similar fashion as a solid lid.

Question 15

Gas-Liquid Equilibrium

Answer 15

As evaporation and condensation proceed, an equilibrium is reached where the rates of the two processes become equal. Once this equilibrium is reached, the pressure that the gas exerts over the liquid is called the vapor pressure of the liquid. Vapor pressure increases as temperature increases since more molecules have sufficient kinetic energy to escape into the gas phase. The temperature at which the vapor pressure of the liquid equals the external pressure is called the boiling point.

Question 16

Liquid-Solid Equilibrium

Answer 16

The liquid and solid phases can also coexist in equilibrium (e.g., the ice water mixture discussed above). Even though the atoms or molecules of a solid are confined to definite locations, each atom or molecules can undergo motions about some equilibrium position. These motions (vibrations) increase when heat is applied. If atoms or molecules in the solid phase absorb enough energy in this fashion, the solid’s 3D structure breaks down, and the liquid phase begins. The transition from solid to liquid is called fusion or melting. The reverse process, from liquid to solid, is called solidification, crystallization, or freezing. The temperature at which these processes occur is called the melting point or freezing point, depending on the direction of the transition. Whereas pure crystals have distinct, very sharp melting points, amorphous solids, such as glass, tend to melt over a larger range of temperature due to their less-ordered molecular distribution.

Question 17

Gas-Solid Equilibrium

Answer 17

A third type of phase equilibrium is that between a gas and a solid. When a solid goes directly into the gas phase, the process is called sublimation. Dry ice (solid CO2) sublimes; the absence of the liquid phase makes it a convenient refrigerant. The reverse transition, from the gaseous to the solid phase, is called deposition.

Question 18

The Gibbs Function

Answer 18

The thermodynamic criterion for each of the above equilibria is that the change in Gibbs free energy must equal zero; deltaG = 0. For an equilibrium between a gas and a solid: deltaG = G(g) - G(s), so G(g) = G(s) at equilibrium The same is true of the Gibbs functions for the other two equilibria.

Question 19

Phase Diagram: Single COmponent

Answer 19

Depicts the phases and phase equilibria of a substance at defined temperatures and pressures. In general, the gas phase is found at high temperature and low pressure; the solid phase is found at low temperature and high pressure; and the liquid phase is found at high temperature and high pressure. The three phases are demarcated by lines indicating the temperatures and pressures at which two phases are in equilibrium. The intersection of the three lines is called the triple point, where all three phases are in equilibrium. The critical point is where the temperature and pressure is where no distinction between liquid and gas is possible.

Question 20

Colligative Properties

Answer 20

Physical properties derived solely from the number of particles present, not the nature of those particles. These properties are usually associated with dilute solutions.

Question 21

Freezing-Point Depression

Answer 21

Pure water (H2O) freezes at 0 degrees Celsius; however, for every mole of solute particles dissolved in 1 L of water, the freezing point, is lowered by 1.86 degrees Celsius. This because the solute particles interfere with the process of crystal formation that occurs during freezing; the solute particles lower the temperature at which the molecules can align themselves into a crystalline structure.

The formula for calculating this freezing-point depression is: ΔT_f = K_fm where ΔT_f is the freezing-point depression, K_f is a proportionality constant characteristic of a particular solvent, and m is the molality of the solution.

Question 22

Boiling-Point Elevation

Answer 22

A liquid boils when its vapor pressure equals the atmospheric pressure. If the vapor pressure of a solution is lower than that of the pure solvent, more energy (and consequently a higher temperature) will be required before its vapor pressure equals atmospheric pressure. The extent to which the boiling point of a solution is raised relative to that of the pure solvent is given by the following formula. deltaTb = Kbm where deltaTb is the boiling-point elevation, Kb is a proportionality constant characteristic of a particular solvent, and m is the molality of the solution. The Kb for water is 0.51 degrees Celsius M^-1.

Question 23

Osmotic Pressure

Answer 23

Consider a container separated into two compartments by a semipermeable membrane (which, by definition, selectively permits the passage of certain molecules). One compartment contains pure water, while the other contains water with dissolved solute. The membrane allows water but not solute to pass through. Because substances tend to flow, or diffuse, from higher to lower concentrations (which increases entropy), water will diffuse from the compartment containing pure water to the compartment containing the water-solute mixture. This net flow will cause the water level in the compartment containing the solution to rise above the level in the compartment containing pure water. Because the solute cannot pass through the membrane, the concentrations of solute in the two compartments can never be equal. However, the pressure exerted by the water level in the solute-containing compartment will eventually oppose the influx of water; thus, the water level will rise only to the point at which it exerts a sufficient pressure to counterbalance the tendency of water to flow across the membrane. This pressure is defined as the osmotic pressure (Pi) of the solution and is given by the formula: Pi = MRT where M is molarity of the solution, R is the ideal gas constant, and T is the temperature on the Kelvin scale. This equation clearly shows that molarity and osmotic pressure are directly proportional (i.e., as the concentration of the solution increases, the osmotic pressure also increases). Thus, the osmotic pressure depends only on the amount of solute, not its identity.

Question 24

Vapor-Pressure Lowering (Raoult’s Law)

Answer 24

When solute B is added to pure solvent A, the vapor pressure of A above the solvent decreases. If the vapor pressure of A above pure solvent A is designated by standardPa and the vapor pressure of A above the solution containing B is Pa, the vapor pressure decreases as follows: deltaP = standardPa - Pa In the late 1800s, the French chemist Francois Marie Raoult determined that this vapor pressure decrease is also equivalent to: deltaP = (Xb)(standardPa) where Xb is the mole fraction of the solute B in solvent A. Becuase Xb = 1 - Xa and deltaP = standardPa - Pa, substituion into the above equation leads to the common form of Raoult’s law: Pa = (Xa)(standardPa) Similarly, the expression for the vapor pressure of the solute in solution (assuming it is volatile) is given by: Pb = (Xb)(standardPb) Raoult’s law holds only when the attraction between molecules of the different components of the mixture is equal to the attraction between the molecules of any one component in its pure state. When this condition does not hold, the relationship between mole fraction and vapor pressure will deviate from Raoult’s law. Solutions that obey Raoult’s Law are called ideal solutions.