Chemistry: The Gas Phase Flashcards
Gas Phase
The atoms or molecules in has move rapidly and are far apart from each other. Only very weak IMF exist between gas particles; this results in certain characteristic physical properties, such as ability to expand to fill any volume and to take on the shape of a container. Furthermore, gases are easily, though not infinitely, compressible.
Gas Pressure Units
1 atm = 760 mm Hg = 760 torr
Standard Temperature and Pressure (STP)
273.15 K and 1 atm. Not to be confused with standard conditions (25 degrees celcius or 298 K).
Boyle’s Law
Experimental studies performed by Robert Boyle in 1660 led to the formulation of Boyle’s law. His work showed that for a given gaseous sample held at constant temperature (isolthermal conditions), the volume of gas is inversely proportional to its pressure PV = K. It’s important to note that the individual values of pressure and volume can vary greatly for a given sample of gas. However, as long as the temperature remains constant and the amount of gas doesn’t change, the product of both P and V will equal the same constant (k). Subsequently, for a given gas under two sets of conditions, the following equation can be derived. P1V1 = k1 = P2V2 or simply P1V1 = P2V2 PV = k or P1V1 = P2V2 where k is a proportionality constant and the subscripts 1 and 2 represent two different sets of conditions.
Law of Charles and Gay-Lussac
Or simply Charles’s law. Was developed during the early 19th century. The law states that at constant pressure, the volume of a gas is directly proportional to its absolute temperature. The absolute temperature is the temperature expressed in Kelvin, which can be calculated from the expression Tk = Tc + 273.15 V/T = k or V1/T1 = V2/T2 where k is a constant and the subscripts 1 and 2 represent 2 different sets of conditions. It’s important to note that the temperature -273.15 is the theoretical lowest attainable temperature, known as absolute zero.
Avogradro’s Principle
In 1811, Amedeo Acogadro proposed that for all gases at a constant temperature and pressure, the volume of the gas will be directly proportional to the number of moles of gas present; therefore, all gases have the same number of moles in the same volume. n/V = k or n1/V1 = n2/V2 The subscripts 1 and 2 once again apply to two different sets of conditions with the same temperature and pressure.
Ideal Gas Law
A theoretical gas whose volume-pressure-temperature behavior can be completely understood by the ideal gas equation is known as an “ideal gas.” The ideal gas law combines the relationships outlined in Boyle’s law, Charles’s law, and Avogadro’s principle to yield an expression that can be used to predict the behavior of a gas. The ideal gas law shows the relationship between four variables that define a sample of gas-pressure (P), volume (V), temperature (T), and number of moles (n). PV = nRT R is the gas constant. Under STP conditions, one mole of gas was shown to have a volume of 22.4 L. Substituting these values into the ideal gas equation gives R = 8.21 x 10^-2 (L x atm)/(mol x K). The gas constant may be expressed in many other units: another common value is 8.314 J/(K x mol), which is derived when SI units of Pascals (for pressure) and cubic meters (for volume) are substituted into the ideal gas law. Note: When calculating using the ideal gas law, it’s important to choose a value of R that matches units of variables.
Density
The mas per unit volume of a substance and, for gases, is usually expressed in units of g/L.
Density: Calculated Using Ideal Gas Law
By rearrangement, the ideal gas equation can be used to calculate the density of gas.
PV = nRT
n = m/MM
PV = (m/MM)RT
d = m/v = {P(MM)}/RT
Density: Calculated Using STP
Another way to find the density of a gas is to start with the volume of a mole of gas at STP, 22.4 L, calculate the effect of pressure and temperature on the volume, and finally calculate the density by dividing the mass by the new volume.
Density: Calculated Using P1V1/T1 = P2V2/T2
The following equation, derived from Boyle’s and Charles’s laws, is used to relate changes in the temperature, volume, and pressure of a gas. P1V1/T1 = P2V2/T2 where the subscripts 1 and 2 refer to the two states of the gas (at STP and under the actual conditions). To calculate a change in volume, the equation is rearranged as follows. V2 = V1(P1/P2)(T1/T2) V2 is then used to find the density of the gas under nonstandard conditions. d = m/V2 Visualizing how the changes in pressure and temperature affect the volume of the gas, can help you check whether or not you have accidentally confused the pressure or temperature value that belongs in the numerator with the one that belongs in the denominator.
Molar Mass
Sometimes the identity of a gas is unknown, and the molar mass must be determined to identify it. Using the equation for density derived from the ideal gas law, the molar mass of a gas can be determined experimentally as follows. The pressure and temperature of a gas contained in a bulb of a given volume are measured, and the weight of the bulb plus sample is found. Then the bulb is evacuated, and the empty bulb is weighed. The weight of the bulb plus sample minus the weight of the bulb yields the weight of the sample. Finally, the density of the sample is determined by divided the weight of the sample by the volume of the bulb. The density at STP is calculated. The molecular weight is then found by multiplying the number of grams per liter by 22.4 liters per mole.
Real Gases
In general, the ideal gas law is a good approximation of the behavior of real gases, but all real gases deviate from ideal gas behavior to some extent, particularly when the gas atoms of molecules are forced into close proximity under high pressure and at low temperature, so that molecular volume and intermolecular attractions become significant.
Deviations Due to Pressure
As the pressure of a gas increases, the particles are pushed closer and closer together. As the condensation pressure for a given temperature is approached, intermolecular attraction forces become more and more significant until the gas condenses into the liquid state. At moderately high pressure (a few hundred atmospheres) the volume of a gas is less than would be predicted by the ideal gas law, due to intermolecular attraction. At extremely high pressure the size of the particles becomes relatively large compared to the distance between them, and this causes the gas to take up a larger volume than would be predicted by the ideal gas law.
Deviations Due to Temperature
As the temperature of a gas is decreased, the average velocity of the gas molecules decreases, and the attractive IMFs become surprisingly significant. As the condensation temperature is approached for a given pressure, intermolecular attractions eventually cause the gas to condense to a liquid state. As the temperature of a gas is reduced toward its condensation point (which is the same as its boiling point), IMA (intermolecular attraction) causes the gas to have a smaller volume than would be predicted by the ideal gas law. The closer the temperature of a gas is to its boiling point, the less ideal is its behavior.