Chapter 5: Gases – Properties, Pressure, and Measurement

Properties of Gases

General Characteristics of Gases

Gases are one of the fundamental states of matter, distinguished by their unique physical properties. Unlike solids and liquids, gases do not have a fixed shape or volume and are highly responsive to changes in pressure and temperature.

• No Definite Volume or Shape: Gases expand to fill the entire volume of their container and assume its shape.

• Compressibility: The volume occupied by a gas changes significantly with pressure, unlike solids and liquids.

• Temperature Dependence: The volume of a gas changes with temperature.

• Miscibility: Gases are easily mixed unless they chemically react with one another.

• Low Density: Gases are typically much less dense than liquids or solids. Gas densities are often expressed in g/L rather than g/mL.

Comparison Table: States of Matter

Property	Solid	Liquid	Gas
Shape	Definite	Indefinite	Indefinite
Volume	Definite	Definite	Indefinite
Compressibility	Very low	Low	High
Density	High	Medium	Low

Characteristic Properties of Gases

Key properties that distinguish gases from other states of matter include:

• Expands to fill its container

• Assumes both the volume and shape of its container

• Readily flows

• Is compressible

• Diffusion within it occurs rapidly

Example: Air in a room will spread out to fill the entire space, regardless of the room's shape or size.

Gas Pressure

Definition and Formula

Pressure is a fundamental property of gases, defined as the force exerted per unit area by gas molecules colliding with the surfaces of their container.

• Formula:

$P = \frac{F}{A}$ Where: $P$ = pressure (Pa, N/m2) $F$ = force (N) $A$ = area (m2)

• Force is calculated as: $F = ma$ (mass × acceleration)

• Units: 1 Pa = 1 N/m2

Example: If the same force is applied over a smaller area, the pressure increases (e.g., pointed heel vs. wide heel).

Calculating Pressure: Real-World Example

Consider a 56 kg person wearing shoes with different heel areas:

• Pointed heel area: 0.45 cm2

• Wide heel area: 16 cm2

Pressure exerted is higher for the pointed heel due to the smaller area.

Comparing Gas Samples by Pressure

Pressure in a gas sample depends on the number of particles and temperature. For samples with the same volume and temperature, more particles mean higher pressure.

• Sample B: Highest pressure (most particles)

• Sample D: Lowest pressure (fewest particles)

• Sample A: Second highest pressure

• Sample C: Second lowest pressure

Atmospheric Pressure

Definition and Origin

Atmospheric pressure is the force exerted by the weight of air in the atmosphere on a unit area at Earth's surface.

• Formula: $P = \frac{F}{A}$

• Standard acceleration due to gravity: $9.81\ \text{m/s}^2$

Example: Atmospheric pressure decreases with altitude because there is less air above the surface to exert force.

Measuring Pressure: Unit Conversion

Barometers and Pressure Units

Pressure is measured using instruments such as barometers. Evangelista Torricelli invented the mercury barometer and defined the unit Torr.

• 1 mmHg = 1 Torr

Common Pressure Units Table:

Pressure Measured in	Unit	Standard Pressure
bar	bar	1 bar
pascal	Pa	100,000 Pa
kilopascal	kPa	100 kPa
Torr (mm Hg)	Torr or mmHg	750.01 Torr
atmosphere	atm	0.98692 atm
pounds per square inch	psi	14.504 psi

Unit Conversion Example: To convert mmHg to bar, use the relationship: $1\ \text{bar} = 750.01\ \text{mmHg}$

Mercury Barometer

A mercury barometer measures atmospheric pressure by the height of a mercury column. The pressure exerted by the atmosphere supports the column of mercury, and the height is proportional to the atmospheric pressure.

• Standard atmospheric pressure: 760 mmHg (1 atm)

Example: If the barometric pressure is 760 mmHg, the height of the mercury column will be 760 mm.

Additional info: These notes cover the introductory concepts of gases, their properties, pressure, and measurement units, as presented in the first part of Chapter 5 of a General Chemistry textbook.