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Chapter Outline

. 10.1 The Properties of Gases . 10.2 Effusion and the Kinetic Molecular Theory of Gases . 10.3 Atmospheric Pressure . 10.4 Relating P,T, and V: The Gas Laws . 10.5 The Combined Gas Law . 10.6 Ideal Gases and the Ideal Gas Law . 10.7 Densities of Gases . 10.8 Gases in Chemical Reactions . 10.9 Mixtures of Gases . 10.10 Solubilities of Gases and Henry’s Law . 10.11 Gas Diffusion: Molecules Moving Rapidly . 10.12 Real Gases

Pressure = force/unit area

Molecules collide with the inside surface of the container. The force of the collision is measured as pressure.

Pressure at Sea Level

Pounds/in2 (psi) 14.7 psi

Atmospheres (atm) 1 atm

Pascals (N/m2) 101.325 X 103 Pa

Torr (mmHg) 760 mmHg

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Torricelli’s Barometer

vacuum The pressure of the Column of atmosphere on the surface mercury of the mercury in the dish is balanced by the 760 mm Hg downward pressure exerted by the mercury in Atmospheric the column. pressure

Elevation and Atmospheric Pressure

0.35 atm

0.62 atm

0.83 atm Sea level

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Chapter Outline

. 10.1 The Properties of Gases . 10.2 Effusion and the Kinetic Molecular Theory of Gases . 10.3 Atmospheric Pressure . 10.4 Relating P,T, and V: The Gas Laws . 10.5 The Combined Gas Law . 10.6 Ideal Gases and the Ideal Gas Law . 10.7 Densities of Gases . 10.8 Gases in Chemical Reactions . 10.9 Mixtures of Gases . 10.10 Solubilities of Gases and Henry’s Law . 10.11 Gas Diffusion: Molecules Moving Rapidly . 10.12 Real Gases

State Variables for a Gas P = pressure V = volume T = temperature n = number of moles

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Boyle’s Law: P and V (n and T held constant) . Gases are compressible • Pressure ↑ as Volume ↓ . Boyle’s Law: • P  1/V (T and n fixed) • or, P × V = constant

• or, P1V1 = P2V2 • Decreasing volume increases number of collisions/area; P↑ (KMT Postulates #3 & 4)

Boyle’s Law and Respiration

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Applying Boyle’s Law Example

A bubble

Explaining Boyle’s Law Using Kinetic Molecular Theory

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Charles’s Law: V and T (n and P held constant) . Charles’s Law: • V  T (P, n constant) V V or, 1 = 2 T T 1 2 Volume of a gas extrapolates to zero at absolute zero (0 K = −273°C). Kinetic energy ↑ as T ↑; force of collisions increases and gas expands to maintain constant P (KMT Post. #3, 4 & 5).

Jacques Alexandre Charles (1796-1823) The French chemist Charles was most famous in his lifetime for his experiments in ballooning. The first such flights were made by the Montgollier brothers in June 1783, using a large spherical balloon made of linen and paper and filled with hot air. In August 1783, however, a different group. supervised by Jacques Charles, tried a different approach. Exploiting his recent discoveries in the study of gases, Charles decided to inflate the balloon with hydrogen gas. Because hydrogen would escape easily from a paper bag, Charles made a bag of silk coaled with a rubber solution. Inflating the bag to its final diameter took several days and required nearly 500 pounds of acid and 1000 pounds of iron to generate the hydrogen gas. A huge crowd watched the ascent on August 27, 1783. The balloon stayed aloft for almost 45 minutes and travelled about 15 miles. When it landed in a village, however, the people were so terrified they tore if to shreds.

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Sample Exercise 10.4: Applying Charles’ Law

Several students at a northern New England campus are hosting a party celebrating the mid-January start of “spring” semester classes. They decide to decorate the front door of their apartment building with party balloons. The air in the inflated balloons is initially 70 oF. After an hour outside, the temperature of the balloons is -12 oF. Assuming no air leaks from the balloons and the pressure in them does not change significantly, how much does their volume change? Express your answer as a percentage of the initial volume.

Explaining Charles’ Law Using Kinetic Molecular Theory

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Avogadro’s Law: V and n (T and P held constant) . Volume is directly proportional to the number of moles of gas, V  n (T, P constant) V  constant n V V or, 1  2 n1 n2 Increasing n increases the number of collisions, gas expands to keep pressure constant (KMT Post. #3 & 4).

Explaining Avogadro’s Law Using Kinetic Molecular Theory

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Amonton’s Law: P and T (n and V held constant) . P  T (n, V constant) P = constant T P P or, 1 = 2 T1 T2 Increasing T will increase force of collisions if volume is kept constant; P will increase (KMT Post. #3, 4 & 5).

Sample Exercise 10.5: Applying Amonton’s Law

Labels on aerosol cans caution against their incineration because the cans may explode when the pressure inside them exceeds 3.00 atm. At what temperature in degrees Celcius might an aerosol can burst if its internal pressure is 2.00 atm at 25 oC?

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Explaining Amonton’s Law Using Kinetic Molecular Theory

Chapter Outline

. 10.1 The Properties of Gases . 10.2 Effusion and the Kinetic Molecular Theory of Gases . 10.3 Atmospheric Pressure . 10.4 Relating P,T, and V: The Gas Laws . 10.5 The Combined Gas Law . 10.6 Ideal Gases and the Ideal Gas Law . 10.7 Densities of Gases . 10.8 Gases in Chemical Reactions . 10.9 Mixtures of Gases . 10.10 Solubilities of Gases and Henry’s Law . 10.11 Gas Diffusion: Molecules Moving Rapidly . 10.12 Real Gases

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The Combined Gas Law

Combining Boyle’s and Charles’ Law (where n is held constant)

Sample Exercise 10.6: Applying the Combined Gas Law

The pressure inside a weather balloon as it is released is 798 mmHg. If the volume and temperature of the balloon are 131 L and 20 oC, what is the volume of the balloon when it reaches an altitude where its internal pressure is 235 mmHg and T = -52 oC?

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