Figure 5.14: A gas whose density is greater than that of air.

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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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Presentation Transcript
Figure 5.15: Finding the vapor density of a substance.

(Example 5.8)

Sample boils

Excess sample escapes

When liquid is gone,

Vapor condenses to solid

Determine mass

Determine density

D= m/v

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

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

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

or

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

Substitute into Ideal Gas Law:

PV = nRT

n = m/Mm

Mm = molecular mass

PV= mRT

Mm

PMm = mRT =dRT

V

A Problem to Consider
• 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.
Stoichiometry Problems Involving Gas Volumes
• 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?
Stoichiometry Problems Involving Gas Volumes
• First we must determine the number of moles of oxygen produced by the reaction.
Stoichiometry Problems Involving Gas Volumes
• Now we can use the ideal gas equation to calculate the volume of oxygen under the conditions given.

O2

Partial Pressures of Gas Mixtures
• 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.
Partial Pressures of Gas Mixtures
• 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.
Partial Pressures of Gas Mixtures
• 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)
A Problem to Consider
• 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
• 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
• Now we can use the ideal gas equation, along with the partial pressure of the hydrogen, to determine its mass.
A Problem to Consider
• From the ideal gas law, PV = nRT, you have
• Next,convert moles of H2 to grams of H2.
Collecting Gases “Over Water”
• 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
• 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.