Chemical Phases and phase equilibria

Question 1

What is the STP of a gas?

Answer 1

The standard temperature and pressure of a gas

Temperature: 0 degrees C, 273.15 K

Pressure: 1 atm, 101.33 kPa, 760 mmHg, 760 torr

Question 2

What is the standard molar volume of a gas?

Answer 2

The volume occupied by one mole of any gas at STP (standard temperature and pressure)

22.4 L/mol

Question 3

What is the kinetic molecular theory of gases? What are its 5 main points?

Answer 3

Describes the behaviour of matter in the gaseous state. A gas that fits this theory exactly is called an ideal gas.

1. Gases are composed of extremely small molecules, separated by distances that are relatively large
2. Molecules of gas are in constant motion, except when they collide with one another
3. Molecules of an ideal gas exert no attractive or repulsive force on one another
4. The collisions experienced by gas molecules do not, on the average, slow them down; rather they cause a change in the direction in which the molecules are moving. If one molecule loses energy as a result of a collision, the energy is gained by the molecule with which is collides. Collisions of the molecules of an ideal gas with the walls of the container result in no loss of energy.
5. The average kinetic energy of the molecules (KE = 1/2 mv^2) increases in direct proportion to the temperature of the gas (KE = 3/2 kT) when the temperature is measured on an absolute scale (ie kelvin) and k is a constant (Boltzmann constant)

Question 4

What four properties of a gas can be experimentally measured?

You need 3 out of these four to calculate the fourth. Present the formula for this (ideal gas law).

Answer 4

1. The weight of a gas, from which the number (N) of molecules of the gas can be calculated
2. The pressure (P) exerted by the gas on the walls of the container in which the gas is placed. (A vacuum is devoid of particles = no pressure)
3. The volume occupied by the gas
4. The temperature (T) of the gas

Ideal gas law: PV = nRT
R = universal gas constant, n = moles of gas

Question 5

Give the ideal gas law and explain its constituents

Answer 5

PV = nRT

R = universal gas constant 0.0821 L-atm/K-mole or 8.31 kPA-dm^3/k-mole

n = moles of gas molecules

Ideal gas law problems can usually be solved just with converting.

1 atm = 760 torr
__ Celsius + 273 = ___ K

Question 6

What does the term partial pressure refer to? What law describes it?

Answer 6

The pressure exerted by one component of the gas mixture if it were occupying the entire volume (V) at the temperature (T)

Dalton’s Law states that the total pressure observed for a mixture of gases is equal to the sum of the pressures that each individual component would exert were it alone in the container

Question 7

How does the behaviour of real gases deviate from ideal gases? (4)

Answer 7

• They do not obey the ideal gas law (and therefore the formula cannot be used to accurately obtain measurements)
• Molecules subject to van der Waal attraction (A real gas acts less like an ideal gas at higher pressures - ideal gas exists when Pressure is Low and Temperature High)
• The particles occupy space, high pressure leads to increase in total volume
• Size and mass of particles effects speed at which they move. Smaller particle size counteracts van der Waal forces, making them faster/more like ideal gas particles. This is shown with Graham’s law, where the rate of movement of a gas (diffusion or effusion) is inversely proportional to the square root of the molecular weight of the gas.

Question 8

What is Graham’s law?

Answer 8

Size and mass of real gas particles effects speed at which they move. Smaller particle size counteracts van der Waal forces, making them faster/more like ideal gas particles. This is shown with Graham’s law, where the rate of movement of a gas (diffusion or effusion) is inversely proportional to the square root of the molecular weight of the gas.

Question 9

How do liquids differ (most noticeably) from gases?

Answer 9

• Viscosity
• Surface tension
• They are Incompressible
• Short distance intermolecular forces hold liquids together (van der Waal forces)

Question 10

List the Van der Waal forces acting on liquids (4)

Answer 10

• Dipole-dipole forces (depend on orientation and distance of molecules, strength diminishes with distance)
• Dipole-induced dipole force: One dipole induces electron distribution in neighbouring molecule that results in attractive force which is inversely proportional to the seventh power of the distance and which is independent of orientation
• London forces: attractive forces acting between nonpolar molecules (due to unsymmetrical instantaneous electron distribution, inducing a dipole in neighbouring molecules)
• Hydrogen bonds: unshielded protons attracted to negative ends of strong dipoles (eg. O, F, N etc.)

Question 11

True or false? A closes system with liquid water and water vapour is homogenous?

Answer 11

False. A homogenous solution has equal physical properties throughout its volume.

Question 12

List the six phase changes

Answer 12

• Sublimation (solid to gas)
• Condensation/deposition (gas to solid)
• Melting
• Freezing
• Condensation (gas to liquid)
• Vaporization/evaporation

These do not involve chemical reactions. Each homogenous phase is separated by other phases by a physical boundary.

Question 13

How can the rate of vaporization be increased? (2)

Answer 13

• Increasing temperature
• Reducing pressure

Both is most effective

Question 14

What is the vapour pressure of a liquid?

Answer 14

When equilibrium is reached in a closed system and the number of molecules in the fixed volume above the liquid remains constant. These molecules exert a constant pressure at a fixed temperature.

Question 15

What is the boiling point of a liquid? How does pressure effect the boiling point? Intermolecular interactions? Molecular weight?

Answer 15

The temperature at which the vapour pressure of the liquid equals the opposing pressure (atmospheric, thus usually air).

Lower pressure produces a boiling point at a lower temperature

Intermolecular interactions (eg. H20, alcohol etc.) decrease vapour pressure and thus raise the boiling point.

Heavier molecules are harder to push into the atmosphere (higher boiling point)

Question 16

What is the freezing point of a liquid? How does atmospheric pressure effect the melting and boiling points?

Answer 16

The temperature at which the vapour pressure of the solid phase equals vapour pressure of the liquid.

Increase in pressure decreases the melting point and increases the boiling point

Question 17

How does the melting point differ for pure and impure solids?

Answer 17

Pure solids completely melt at one temperature (melting point). Impure solids begin to melt at one temperature and finish completely melting at another.

Question 18

What is the triple point on a phase diagram? What is the metastable equilibrium?

Answer 18

• The triple point is the pressure/temperature where solid, liquid and vapour are in equilibrium. Ice and water have the same fixed vapour pressure at this temperature.
• A state of pseudo-equilibrium having higher free energy than the true equilibrium state. Eg. supercooled water and its vapour

Question 19

If presented with four chemical compounds, the molecular weight of each and the melting point of each, how can you determine which two best represent the effect of molecular weight on melting point?

Answer 19

The two compounds which are most structurally similar should be compared to assess the effect that molecular weight has on melting point (or any other attribute).

You do not necessarily need to pick out the two compounds with the most varying molecular weights, just the most similar structures.

Question 20

A student prepared a .1 M aqueous solution of crotonic acid and a .1 M aqueous solution of oxalic acid, then adjusted the pH of each to 4.7 by adding NaOH. Which solution has a lower freezing point?

Crotonic acid pKa: 4.69

Oxalic acid pKa: 3.14, 4.77

Answer 20

The oxalic acid solution would have a lower freezing point because it contains a greater molar concentration of solute particles.

The freezing point depression of an aqueous solution is a colligative property. Given two solutions, the one with the greater number of solute particles per liter of solution freezes at the lower temperature.

Oxalic acid is diprotic and ionizes in accord with the pKa values shown to a greater extent than does crotonic acid. Subsequently, oxalic acid requires more NaOH than does crotonic acid to reach a pH of 4.7, and oxalic acid produces a larger number of particles in solution.

Question 21

A student has a thin copper beaker containing 100g of a pure metal in the solid state. The metal is at 215 degrees celsius, its exact melting temperature. If the student lights a Bunsen burner and holds it for a fraction of a second under the beaker, what will happen to the metal?

Answer 21

A small amount of the metal will turn to liquid, with the temperature remaining the same.

Melting occurs at a constant temperature because a certain amount of energy, the latent heat of fusion, is needed to convert a substance from its solid to liquid state. The temperature of the metal will not increase above its melting point until all the metal has melted. The small amount of heat supplied by the Bunsen burner is insufficient to melt 100 g of the metal, but it could melt a small amount at the constant temp of the melting point.

Question 22

Incandescent light bulbs have a wire filament that channels a current through it. Heating the filament produces light, but also causes the vaporization of the filament.

Why does an incandescent light have a longer life when an inert gas is used, rather than a vacuum?

Answer 22

IF the filament is in a vacuum, it will vaporize at a faster rate than it would at normal atmospheric pressure (remember phase diagram).

As pressure decreases at a specific temperature, it becomes possible for a solid to sublimate directly to a gas (vaporization). A vacuum was initially used to prevent oxygen from oxidizing the filament, an inert gas is better because it does not require a vacuum.