Study Notes: Chapter 11 – Gases (Introductory Chemistry)

Gases: Properties, Laws, and Applications

Introduction to Gases

Gases are one of the fundamental states of matter, characterized by their ability to expand and fill any container. Understanding the behavior of gases is essential in chemistry, as it explains phenomena from breathing to weather patterns.

Extra-Long Straws: Atmospheric Pressure and Gas Behavior

How Straws Work

• Pressure Difference: Drinking from a straw creates a pressure difference between the inside and outside of the straw, causing liquid to be pushed up.

• Atmospheric Molecules: The pushing force is due to atmospheric molecules (mainly nitrogen and oxygen) colliding with surfaces.

• Maximum Straw Length: Even with perfect materials and a vacuum, atmospheric pressure can only push liquid up to about 10.3 m (34 ft) in a straw. This is because a column of water (or soda) 10.3 m high exerts the same pressure as the atmosphere at sea level.

Example: If you try to drink soda through a straw longer than 10.3 m, the atmospheric pressure is insufficient to push the liquid up the straw.

Kinetic Molecular Theory: A Model for Gases

Fundamental Postulates

• Constant, Straight-Line Motion: Gas particles move in constant, straight lines.

• No Interactions: Gas particles do not attract or repel each other; collisions are elastic, like billiard balls.

• Large Spaces: There is a lot of empty space between gas particles compared to their size.

• Kinetic Energy and Temperature: The average kinetic energy of gas particles is proportional to the temperature in kelvin. As temperature increases, particles move faster.

Additional info: The kinetic molecular theory explains why gases are compressible, expand to fill containers, and have low densities.

Properties of Gases Explained by Kinetic Molecular Theory

• Compressibility: Gases are compressible due to the large amount of empty space between particles.

• Shape and Volume: Gases assume the shape and volume of their container because intermolecular attractions are negligible.

• Low Density: Gases have much lower densities than liquids and solids.

Example: 350 mL of liquid water converts to 595 L of steam at 100°C and 1 atm, showing the vast increase in volume and decrease in density.

Pressure: The Result of Constant Molecular Collisions

Definition and Effects

• Pressure: The force per unit area resulting from collisions of gas particles with surfaces.

• Formula: $\text{pressure} = \frac{\text{force}}{\text{area}}$

• Factors Affecting Pressure: Number of gas particles in a given volume; more particles mean higher pressure.

• Real-World Effects: Pressure allows us to drink from straws, inflate objects, and breathe. Atmospheric pressure changes with altitude, affecting the body (e.g., ear pain in airplanes).

Units of Pressure

Common Pressure Units

• Atmosphere (atm): Average pressure at sea level.

• Pascal (Pa): SI unit, defined as 1 newton per square meter. $1\,\text{Pa} = 1\,\text{N}/\text{m}^2$

• Millimeter of Mercury (mm Hg): Based on barometer measurements; 1 atm = 760 mm Hg.

• Torr: 1 mm Hg = 1 torr.

• Other Units: Inches of mercury (in. Hg), pounds per square inch (psi).

Unit	Average Air Pressure at Sea Level
Pascal (Pa)	101,325 Pa
Atmosphere (atm)	1 atm
Millimeter of mercury (mm Hg)	760 mm Hg
Torr (torr)	760 torr
Pounds per square inch (psi)	14.7 psi
Inches of mercury (in. Hg)	29.92 in. Hg

Pressure Unit Conversion Example

• To convert 0.311 atm to mm Hg: $0.311\,\text{atm} \times \frac{760\,\text{mm Hg}}{1\,\text{atm}} = 236\,\text{mm Hg}$

Everyday Chemistry: Airplane Cabin Pressurization

Effects of Low Pressure

• At high altitudes (25,000–40,000 ft), atmospheric pressure is below 0.50 atm, much lower than at sea level.

• Low pressure leads to reduced oxygen levels, causing dizziness, headache, and even unconsciousness.

• Cabin air is pressurized using compressed air from jet engines, mixed and regulated to maintain safe pressure.

• Federal regulations require cabin pressure to be greater than the equivalent of outside air pressure at 8,000 ft (about 0.72 atm).

Example: Oxygen masks are used if cabin pressurization fails.

Boyle's Law: Pressure and Volume

Relationship Between Pressure and Volume

• Inverse Proportionality: At constant temperature and amount of gas, pressure increases as volume decreases, and vice versa.

• Mathematical Expression: $P \propto \frac{1}{V}$ $PV = \text{constant}$

• Application: Used in hand pumps and scuba diving to explain changes in pressure and volume.

Example: If a diver ascends rapidly from depth, the pressure in the lungs drops, causing the volume of air to expand and potentially damage the lungs.