Gases Liquids and Solids

Kinetic Molecular Theory of Matter

1) All matter is composed of small particles

2) The particles are in constant motion and therefore possess kinetic energy

3) The particles can have attractions and repulsions between them

4) The kinetic energy of the particles increases as temperature increases

5) The particles collide elastically and can transfer energy to each other during these collisions

Consider the phases of matter in light of the kinetic molecular theory

(kinetic energy verses attractions)

Pressure: a measurement in the gas phase

Pressure is the force applied to a given areaPressure is the force applied to a given area

P = F / AP = F / A

Pressure can increase by increasing the force or Pressure can increase by increasing the force or decreasing the areadecreasing the area

Units: Units: mm Hgmm Hgatmospheres (atm)atmospheres (atm)torrtorr

Units of Pressure

1 mm Hg / 1 torr1 mm Hg / 1 torr

1 atm / 760 mm Hg1 atm / 760 mm Hg

1 atm / 760 torr1 atm / 760 torr

1 atm is a standard 1 atm is a standard unit of pressure and unit of pressure and corresponds to corresponds to pressure at sea levelpressure at sea level

A Barometer

Boyle’s Law

At constant temperature

As volume goes down, pressure goes up (Think about squeezing a balloon)

P1 x V1 = P2 x V2

Pressure and volume can have any units as long as they are the same on both sides

Boyle’s Law in Practice

Filling a syringe Increasing the volume of the syringe decreases its pressure so the liquid flows in

This is how inhaling works too.Expanding your lungs creates more volume and less pressure so air from the room rushes in

Charles’s Law

At constant pressure

As temperature goes down, volume goes down (Think about tires in the winter and summer)

Volume can have any units as long as they are the same on both sides.

Temperature must be Kelvin (K = C + 273)

V1 V2

T1 T2

=

Combined Gas Law

If you know this… you know both Boyle’s and Charles’s

These types of problems give one set of conditions… (V1, T1 etc….)

then a new set of conditions…. (V2, T2 etc….)

And you have one single variable to solve for

P1 V1 P2 V2

T1 T2

=

Try some

A gas has a volume of 250.0 mL at 1.75 atm of pressure. If the pressure is changed to 765 mm Hg, what is the new volume in mL?

A gas occupies 500.0 L at 1.00 atm when the temperature is 61 C. In August the temperature climbs to 99 C and the gas is in the same container with the same volume. What is it’s new pressure?

Ideal Gases

Ideal gases: Have perfectly elastic collisions with no attractions at all when collidingOccupy no volume compared to the volume of the container

Real gases however do have volumes and do have attractions

When are real gases like ideal gases? Under conditions of large container volume and/or high temperature.

Ideal Gas Law

PV = nRT

P - pressure in atmV – volume in Ln – number of moles of gasT – temperature in KelvinR – ideal gas constant

0.0821 L ∙ atm / mol ∙ K

These problems have only one set of conditions and you solve for one variable

Keep in mind…..

Pressure – may need to convert mm Hg or torr into atm

Volume – may need to convert into liters

n – may need to convert grams into moles

Temperature – may need to convert celsius into Kelvin

Try this….

How many liters would be occupied by 12.05 moles of H2 gas that is at 125° C and exerts 775 mmHg of pressure?

Daltons Law of Partial Pressures

In a mixture of gases, the total pressure is the sum of the pressures of the individual gases

PT = P1 + P2 + P3 …..

Phase Changes

The boiling process

Volatility – how easily something evaporates

Boiling point – the temperature at which a liquid boils

Vapor pressure – the pressure exerted by a vapor as it evaporates from a liquid

Boiling occurs when:

vapor pressure = surrounding pressure

Boiling point of water at different locations (different pressures)

Intermolecular forces

Ion – Ion Strongest

Ion - Dipole

Dipole - Dipole

London dispersion Weakest

What is present in each of these?