chapter 13 Essay

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CHAPTER 13: STATES OF MATTER

GASES

Properties of Gases:
• Particles spread far apart
• Dispersion forces between molecules (some dipole)
• Exert pressure

Kinetic Theory: (for gases only)

The particles in a gas are assumed to be small, hard spheres with an insignificant volume

Particles are far apart with no attraction between particles (empty space)

The motion of particles in a gas is rapid and random.

Gases fill containers, particles travel in straight paths until they collide with walls or other particles

All collisions between gas particles are perfectly elastic.

Kinetic energy is transferred without any being lost because they have no attraction to each other

Pressure:

Definition:

Force exerted by a gas per unit surface area on an object

What creates pressure?

Result of rapidly moving particles in a gas simultaneously colliding with the surface or container walls (more hits = more pressure; less hits = less pressure)

Atmospheric Pressure: Barometer

Number of collisions of atoms and molecules in air with the surface (measured by barometer)

Pressure Units: 1 atmosphere = 760 mmHg = 101.3kPa

Make the following pressure conversions

4.3 atm to mmHg

4.3 atm 760 mmHg = 3268 mmHg 1 atm

2.25atm to kPa

2.25 atm 101.3 kPa = 228 kPa 1 atm

450mmHg to kPa

450 mmHg 101.3 kPa = 60 kPa 760 mmHg

DEMONSTRATIONS:

Coke Can
Heating the can causes air particles to move faster  more collisions  increase in pressure (but because the top of the can is open, particles can leave allowing the pressure inside and outside the can to be equal)
Placing the can into ice water cools its down quickly and the molecules can’t move as much so the pressure outside the can is higher than inside the can which causes the air pressure to crush the can
Balloon in Bottle

Water bottle with card
Air pressure > water pressure so the card sticks to the opening of the bottle rather than the water pushing the card away

Volume:

Definition: measure of space occupied by matter

Units: 1000 mL = 1 L = 1 dm3 1 mL = 1 cm3

Kinetic Energy and Temperature:

Kinetic Energy: energy of motion – dependent on mass (how heavy) and velocity (speed)

Temperature: average kinetic energy of particles in a sample (think speed)

a. Which point on each curve represents the average kinetic energy? (label on diagram)
Top of each peak
b. Which point on each curve represents the slowest moving particles? (label on diagram)
@ 0 m/sec
c. Which point on each curve represents the fastest moving particles? (label on diagram)
@ 1200 m/sec for 25°C and 2000 m/sec for 1000°C
d. What do you think would happen to the first shape of the curve if the temperature were even lower?
More extreme curve (higher peak)

Kelvin Temperature:

Units: 1 K = 1 °C K = °C +273

Make the following temperature conversions:

100°C to K 323 K to °C
100 + 273 = 373 K 323 – 273 = 50 °C

Why Kelvin Temperature?
Kelvin temperature is directly proportional to the average kinetic energy of the particles of the substance

What is Absolute Zero?
All
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