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

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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

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.

If we solve this equation for the molecular mass, we obtain

Molecular Weight Determination- If we substitute this in the ideal gas equation, we obtain

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)

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

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

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.

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

Return to Slide 44

Figure 5.26: Model of gaseous effusion.

Return to Slide 45

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