Unit 10.1

Pressure and temperature

Pressure

Pressure = Force/Area (N/m2= pascal)

-more force over same area = higher pressure

-more area with the same force = lower pressure

Atmospheric pressure=pressure of air in the Earth’s atmosphere.

Atmospheric pressure decreases as altitude increases.

Measuring Pressure

The first device for measuring atmosphericpressure was developed by Evangelista Torricelliduring the 17th century.

The device was called a “barometer”

Baro = weightMeter = measure

An Early Barometer

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

Standard Pressure

atomospheric pressure at sea level

1 standard atmosphere (atm) 101.3 kPa (kilopascals)- SI unit 14.7 lbs/in2

760 mm Hg (millimeters of mercury) 760 torr

Use standard pressure to convert between pressure units

The Kelvin Scale

Converting Celsius to Kelvin

Kelvins = C + 273

°C = Kelvins - 273

Unit 10.2

Properties of gases, liquids, and solids

Class starter

Covert the following

1. 56.1 torr to kPa

2. 780.0 torr to atm

3. 2.30 atm to mmHg

4. 45 ْC to Kelvin

5. 560 K to ْC

Kinetic Molecular Theory

1. pertaining to motion.

2. caused by motion.

3.characterized by movement: Running and dancing are kinetic activities.

ki⋅net⋅ic

Origin: 1850–55; < Gk kīnētikós moving, equiv. to kīnē- (verbid s. of kīneîn to move) + -tikos

Source: Websters Dictionary

Kinetic Molecular Theory of gases

Particles of matter are ALWAYS in motion

Volume of individual particles is zero.

Collisions of particles with container walls cause the pressure exerted by gas.

Particles exert no forces on each other.

Average kinetic energy is proportional to Kelvin temperature of a gas.

Properties of Gases

Gases expand to fill their containers

Gases are fluid – they flow

Gases have low density

1/1000 the density of the equivalent liquid or solid

Gases are compressible

Gases effuse and diffuse

Gases exert pressure on the walls of their containers

Kinetic Energy of Gas Particles

At the same conditions of temperature, all gases have the same average kinetic energy.

m = mass

v = velocity

2

2

1mvKE

At the same temperature, light moleculesmove FASTER than heavy molecules

Diffusion describes the mixing of gases. The rate of diffusion is the rate of gas mixing.

Diffusion is the result of random movement of gas molecules

The rate of diffusion increases with temperature

Small molecules diffuse faster than large molecules

Diffusion

Properties of liquids

Particles are closer together than gases and have more effective intermolecular forces than gases

lower kinetic energy than gases

High density

Relatively incompressible

Properties of liquids

• Surface tension- inward pull of molecules at the surface of a liquid which gives the liquid a “skin”

-liquids with stronger intermolecular forces have higher surface tension

Capillary action- attraction of the surface of a liquid to the surface of a solid. This attraction can pull a liquid against gravity.

Example: water moving up a paper towel, meniscus

chromatography uses capillary action to separate things

• Volatile liquids evaporate readily

• Nonvolatile liquids do not evaporate readily

• The weaker the intermolecular forces, the more volatile the liquid (easier to escape the surface of the liquid and evaporate)

Properties of solids

• Particles are closely packed with very effective intermolecular forces

• Very low kinetic energy

• High density – in general substances are most dense as a solid (exception: water )

• Relatively incompressible

Unit 10.3

Phase changes and Heating curves

Class starter 1. Which of the following will diffuse the fastest?

Ne at 20 ⁰C or Ne at 40 ⁰C

2. Which will diffuse the fastest?

Cl2 at 20 ⁰C or F2 at 20 ⁰C

3. In each pair, which molecule has the highest kinetic energy?

a. Ar at 30 ⁰C or Ne at 30 ⁰C

b. Ar at 30 ⁰C or Ar at 40 ⁰C

4. 40 ⁰C = _____K

5. Which liquid is more volatile? NH3 or CH4

Phase changes

Phase changes• Vaporization- liquid changes to a gas

– Evaporation: particles escape from the surface of a nonboiling liquid

– Boiling: change of a liquid to bubbles of vapor

Energy change in phase changes

• Exothermic: heat is released

- Freezing, condensation, deposition

– Feels warm around it, temperature increases

• Endothermic: heat is absorbed

– Melting, boiling, sublimation

– Feels cold around it, temperature decreases

kinetic energy tells how fast molecules are moving. (proportional to temperature)

Potential energy tells the potential the atoms have to bond.

-decrease potential energy - bonds are formed.

-increase potential energy – bond are broken

Heating curve

Heating curve explanations

• A positive slope shows that the temperature is increasing (Kinetic energy is increasing). Molecules are speeding up.

• When a substance melts or boils the temperature remains constant (kinetic energy stays the same) because all of the energy goes into breaking the intermolecular forces instead of speeding the molecules up. (potential energy increases)

Heating curve explanations

Molar heat of fusion ( ∆Hfus) : amount of energy needed to melt one mole of a solid

Molar heat of vaporization (∆Hvap) : amount of energy needed to vaporize one mole of a liquid

∆Hvap > ∆Hfus

Why? When vaporizing a substance all bonds need to be

broken, when melting a substance only some bonds are broken, most remain intact.

• Which of the following would you predict to have a higher heat of vaporization?

CH4 (nonpolar) or H2O (polar)

• H2O has stronger intermolecular forces so it would require more energy to pull them apart.

• Stronger intermolecular forces means higher heats of fusion and vaporization