States of Matter Kinetic-Molecular Theory – describes the behavior of gases Three assumptions: 1)All gases consist of small particles 2)Constant random.

States of MatterKinetic-Molecular Theory – describes the behavior of gases

Three assumptions:

1)All gases consist of small particles

2)Constant random motion experiencing elastic collisions

3)Kinetic energy depends on mass and velocity. At a given temperature, all gases have the same kinetic energy.

K.E. = ½ mv2

m mass in kg

v velocity m/s

K. E. Proportional to temperature

Consider Nitrogen and oxygen at room temp.

Which one has higher K.E.?

Which one has faster moving molecules?

Behavior of gases

• Low density

• Compression and expansion

• Diffusion-movement of one material through another

• Effusion-gas escaping through a small hole

• Graham’s law

Consider gas 1 and gas 2 at room temp.

• Do page 388 1-3

Gas Pressure

• Pressure – force per unit area

• Barometer – measures atm pressure

• At sea level at 0 degrees C, average air pressure is 760. mm Hg. This is the same as 1.00 atm or 760. torr or 101.3 kPa

• Be able to do conversions

• Manometer – measures gas pressure in a closed container (see page 389)

• Open manometers

Dalton’s law of partial pressures – the individual pressures add up to

the total

• Do page 392 4-6

Forces of AttractionIntermolecular forces

• Dispersion forces - weak forces due to a temporary induced dipole

• Exist between all particles but only significant if they are the only forces present

• Explains why Cl2 and F2 are gases but Br2 is a liquid and I2 is a solid

• Dipole-dipole• Hydrogen bonds

Liquids

• Density and compression

• Fluidity

• Viscosity – resistance to flow, decreases with temperature

• Surface tension

• Surfactants

• Capillary action

Solids

• Density

• Crystalline

• Simple cubic cell

• Body-centered cubic cell

• Face-centered cubic cell

• Molecular solids

• Covalent network solids (macromolecular)

• Ionic solids

• Metallic solids

• Amorphous solids

Phase changes

• Melting • Vaporization• Vapor pressure • Boiling – know this definition• Sublimation• Condensation• Deposition• freezing

Phase Diagrams

q = mass X change in temp X specific heat

q = mass X enthalpy

For water: specific heat of steam is 2.02 J/g C

specific heat of liquid is 4.18 J/g C

specific heat of ice is 2.06 J/g C

Enthalpy of fusion is 334 J/g

Enthalpy of vaporization is 2260 J/g

Energy and change in state

• How much energy is needed to convert 10.0 g of ice at -15.0 degrees C to steam at 130.0 degrees C?

Energy and change in state

• How much energy is needed to melt 25.4 grams of iodine? The enthalpy of fusion of iodine is 61.7 J/g.

• How much energy is needed to melt 4.24 grams of Pd? The enthalpy of fusion of Pd is 162 J/g.

• Find the heat needed to raise the temperature of 5.58 kg of iron from 22.0 degrees C to 1000.0 degrees C.

• Find the energy change when 28.9 g of copper is cooled from it’s melting point to 25.0 degrees C. Is this energy gained or released and how can you tell?

Coffee cup calorimeter

example

• A 2.35 g sample of unknown metal at 100.0 degrees C is placed in 50.0 ml of water originally at 23.0 degrees C. The maximum temperature of the water is 23.3 degrees C. What is the specific heat of the unknown metal?

Ice lab

• Design a lab that allows you to find the molar heat of fusion of ice.