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# Figure 5.14: A gas whose density is greater than that of air. - PowerPoint PPT Presentation

Figure 5.14: A gas whose density is greater than that of air. Figure 5.15: Finding the vapor density of a substance. (Example 5.8). Sample boils Excess sample escapes When liquid is gone, submerge flask in ice water, Vapor condenses to solid Determine mass Determine density. D= m/v.

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(Example 5.8)

Sample boils

Excess sample escapes

When liquid is gone,

Vapor condenses to solid

Determine mass

Determine density

D= m/v

• If we look again at our derivation of the molecular mass equation,

we can solve for m/V, which represents density.

• Calculate the density of ozone, O3 (Mm = 48.0g/mol), at 50 oC and 1.75 atm of pressure.

or air.

Molecular Weight Determination

• In Chapter 3 we showed the relationship between moles and mass.

PV = nRT

n = m/Mm

Mm = molecular mass

PV= mRT

Mm

PMm = mRT =dRT

V

Molecular Weight Determination

• If we substitute this in the ideal gas equation, we obtain

• A 15.5 gram sample of an unknown gas occupied a volume of 5.75 L at 25 oC and a pressure of 1.08 atm. Calculate its molecular mass.

• Consider the following reaction, which is often used to generate small quantities of oxygen.

• Suppose you heat 0.0100 mol of potassium chlorate, KClO3, in a test tube. How many liters of oxygen can you produce at 298 K and 1.02 atm?

• First we must determine the number of moles of oxygen produced by the reaction.

• Now we can use the ideal gas equation to calculate the volume of oxygen under the conditions given.

O2

• The composition of a gas mixture is often described in terms of its mole fraction.

• Themole fraction, c , of a component gas is the fraction of moles of that component in the total moles of gas mixture.

• The partial pressure of a component gas, “A”, is then defined as

• Applying this concept to the ideal gas equation, we find that each gas can be treated independently.

• Dalton’s Law of Partial Pressures: the sum of all the pressures of all the different gases in a mixture equals the total pressure of the mixture. (Figure 5.16)

Figure 5.17: An illustration pressures before mixing.of Dalton’s law of partial pressures after mixing.

A Problem to Consider pressures before mixing.

• Given a mixture of gases in the atmosphere at 760 torr, what is the partial pressure of N2 (c = 0 .7808) at 25 oC?

A Problem to Consider pressures before mixing.

• Suppose a 156 mL sample of H2 gas was collected over water at 19 oC and 769 mm Hg. What is the mass of H2 collected?

• Table 5.6 lists the vapor pressure of water at 19 oC as 16.5 mm Hg.

A Problem to Consider pressures before mixing.

• Now we can use the ideal gas equation, along with the partial pressure of the hydrogen, to determine its mass.

A Problem to Consider pressures before mixing.

• From the ideal gas law, PV = nRT, you have

• Next,convert moles of H2 to grams of H2.

Collecting Gases “Over Water” pressures before mixing.

• A useful application of partial pressures arises when you collect gases over water. (see Figure 5.17)

• As gas bubbles through the water, the gas becomes saturated with water vapor.

• The partial pressure of the water in this “mixture” depends only on the temperature. (see Table 5.6)

A Problem to Consider pressures before mixing.

• Suppose a 156 mL sample of H2 gas was collected over water at 19 oC and 769 mm Hg. What is the mass of H2 collected?

• First, we must find the partial pressure of the dry H2.

Conceptual Problem 5.27 pressures before mixing.

Figure 5.27: The hydrogen fountain. pressures before mixing.Photo courtesy of American Color.