1. Gases
Chapter 13-1

2. The Nature of Gases Kinetic refers to motion
The energy an object has because of it’s motion is called kinetic energy
The kinetic –molecular theory states that the tiny particles in all forms of matter are in constant motion!

3. The Nature of Gases
Three basic assumptions of the kinetic theory as it applies to gases:
#1. Gas is composed of particles- usually molecules or atoms
Small, hard spheres
Insignificant volume; relatively far apart from each other
No attraction or repulsion between particles

4. The Nature of Gases
#2. Particles in a gas move rapidly in constant random motion
Move in straight paths, changing direction only when colliding with one another or other objects
Average speed of O2 in air at 20 oC is an amazing 1700 km/h (1062 mph)!

5. The Nature of Gases
#3. Collisions are perfectly elastic- meaning kinetic energy is transferred without loss from one particle to another- the total kinetic energy remains constant

6. The Nature of Gases Two factors that determine kinetic energy is:
Mass and Velocity
KE = 1/2mv2
KE and Temperature are also related. At a given temp, ALL gases have the same average KE

7. Behavior of Gases
Low Density – gas is less dense (mass per volume) than liquid or solid
Compression/Expansion – as the volume is compressed, the density increases; as volume expands, the density decreases

8. Behavior of Gases Diffusion/Effusion
Diffusion is when molecules disperse through the air (perfume from one room to another)
Effusion is when gas escapes through a tiny hole (as in a balloon)
The rate of effusion is inversely proportional to the square root of its molar mass
Rate of effusion =
𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠

9. Gas Pressure
Gas Pressure – defined as the force exerted by a gas per unit surface area of an object
Due to: a) force of collisions, and b) number of collisions
No particles present? Then there cannot be any collisions, and thus no pressure – called a vacuum

10. Gas Pressure
Atmospheric pressure results from the collisions of air molecules with objects
Decreases as you climb a mountain because the air layer thins out as elevation increases
Barometer is the measuring device for atmospheric pressure, which is dependent upon weather & altitude

11. Measuring Pressure The first device for measuring atmospheric
pressure was developed by Evangelista Torricelli
during the 17th century.
The device was called a “barometer”
Baro = weight
Meter = measure
Torricelli

12. Gas Pressure The SI unit of pressure is the pascal (Pa)
At sea level, atmospheric pressure is about kilopascals (kPa) (which also = 1atm)
Older units of pressure include millimeters of mercury (mm Hg), and atmospheres (atm) – both of which came from using a mercury barometer

13. Gas Pressure
Mercury Barometer – Fig. 13.2, page 386 – a straight glass tube filled with Hg, and closed at one end; placed in a dish of Hg, with the open end below the surface
At sea level, the mercury would rise to 760 mm high at 25 oC- called one standard atmosphere (atm)

14. An Early Barometer
The normal pressure due to the atmosphere at sea level can support a column of mercury that is 760 mm high.

15. Gas Pressure This is called Standard Temperature and Pressure, or STP
For gases, it is important to relate measured values to standards
Standard values are defined as a temperature of 0 oC and a pressure of kPa, or 1 atm
This is called Standard Temperature and Pressure, or STP

16. Law of Partial Pressure
Dalton’s Law of Partial Pressure states that the total pressure of a mixture of gases is equal to the sum of the pressures of all the gases in the mixture
Depends on: number of moles of gas, the size of the container and temp.
Ptotal = P1 + P2 + P3

17. Practice Problem
A mixture of oxygen (O2), carbon dioxide (CO2) and nitrogen (N2) has a total pressure of 0.97 atm. What is the partial pressure of O2, if the partial pressure of CO2 is 0.70 atm and the partial pressure of N2 is 0.12 atm?