Chapter 10: States of Matter

10.1 The Nature of Gases

Kinetic Theory

Gases are made up of atoms or molecules. The atoms can be modeled as small, hard spheres like marbles. They are far apart from each other. They occupy insignificant volume (relative to the amount of empty space between them). No significant attractive or repulsive forces exist. Particles move in random, constant motion, in straight paths between collisions.

When they collide, it is a perfectly elastic collision.

Pressure of Gases

Can you name a unit of pressure?

psi = pounds per square inch =

Gas pressure is the force per unit surface area (of the container).

What is creating the force by a gas in a container?

When the gas molecules collide with the wall of the container, it creates a force against the wall of the container. Pressure is that force per unit area.

If the molecules collide with the wall when they are going FASTER, is that more pressure or less pressure?

If there are more collisions at the same speed, does that make the pressure more or less?

Units of pressure

101.3 kPa = 1 atm = 760 mm Hg = 14.7 psi (kiloPascals) (atmospheres) (mm Mercury)(lbs/in2) Where does the 1 atmosphere come from? Earth’s atmosphere is forced downwards towards the

center of the Earth by Earth’s gravity. It is the weight of the air column above a unit area on the surface of the Earth that causes “atmospheric pressure”.

Atmospheric pressure is defined to be 1 atm at sea level.

Up at high altitudes, the air column is shorter, so the atmospheric pressure is lower up there (and the air gets less debse or “thin” up there).

Barometers are used to measure atmospheric pressure.

Atmospheric pressure at sea level

At sea level, atmospheric pressure is 1 atm or 101.3 kPa or 760 mm Hg, as measured by a barometer.

In the drawing at the left, the pressure at sea level forces mercury up the tube to a height of 760 mm.

That is why 760 mm Hg = 1 atm

Conversions between pressure units

Example: Convert 5.2 atm to kPa

Example: Convert 201 kPa to mm Hg

Kinetic Energy and Kelvin Temperature

The hotter the temperature, the higher the average kinetic energy (the faster the particles are moving).

At 0K, the motion stops, as you can see below:

Simulation showing KE vs. temperature: states-of-matter-basics_en.jar

Kinetic Energy and Temperature Temperature is related to kinetic energy.

Temperature on the Kelvin scale is directly proportional to the average kinetic energy.

Kinetic energy KE = ½ m v2 so kinetic energy is proportional to velocity of the particle squared.

When you heat gas particles, they move around faster, more KE.

The average KE of the sample is proportional to temperature in Kelvin

KE and Temperature

Example 1: What happens to KE if you triple the temperature in Kelvin?

KE triples

Example 2: What happens to KE when you go from -123oC to 27oC?

That temperature change is going from 150K to 300K. Does that help?

KE doubles

Also, just as a reminder, STP (Standard temperature and pressure) are as follows:

0o C and 1 atm (or you can use 101.3 kPa)

Section 10.2 Liquids

Unlike gases, liquids are held together by attractive forces.

Liquids have fixed volumes, meaning they are not easily compressed.

Liquids take the shape of the container.

Liquid particles are free to slide past one another.

The conversion of a liquid to a gas or vapor is called vaporization.

When the vaporization takes place at the surface of a liquid that is not boiling, it is called evaporation.

evaporation

Liquid Gas

condensation

It takes energy for a liquid particle to overcome the attractive forces in the liquid phase and evaporate to become a gas.

Only molecules that have enough energy evaporate (this happens on the surface of the liquid).

Liquids, continued In an open container, water

molecules evaporate from the liquid (and notice how it also shows some condensing). The molecules that evaporate escape the system, so equilibrium is never reached.

If it is a closed vessel, the water vapor above the liquid reaches a dynamic equilibrium with the liquid phase, where the net evaporation goes to zero.

Gas/vapor equilibrium

With the closed container, the equilibrium between liquid and gas/vapor particles is like this:

evaporation

Liquid Gas (vapor)

condensation

The definition of equilibrium is when the

Rate of evaporation = Rate of condensation

So there is no longer any net evaporation, and the liquid level stays constant.

Liquids

When evaporation occurs, say after you have jumped out of the swimming pool. What effect does that have on you? The water molecule takes heat from your body and

applies it towards making the phase change from liquid to gas.

When it escapes to the gas phase and departs, it takes some of your heat with it.

Evaporation is a cooling process.

What is vapor pressure? Vapor pressure only applies to sealed containers

where the gas is trapped.

The gas molecules collide with the walls of the container. Each collision is a small force against the wall of the container. If you add up all these small forces and divide by the surface area of the container:

Pressure = Force / Area

This is the vapor pressure- the pressure applied to the walls by the force of the vapor.

Higher temperature >> more evaporation >> more vapor >> higher vapor pressure.

Vapor pressure is measured with a manometer (page 277)

Vapor PressureVapor Pressures (kPa) of Several Substances at Various Temperatures

0o C 20o C 40o C 60o C 80o C 100o C

Water 0.61 2.33 7.37 19.92 47.34 101.33

Ethanol 1.63 5.85 18.04 47.02 108.34 225.75

Diethyl ether

24.70 58.96 122.80 230.65 399.11 647.87

What does that mean when it says that the vapor pressure of water at 100o C is 101.3 kPa? (recall that 101.3 kPa = 1 atm)

Would you say ethanol and diethyl ether are more or less volatile than water?

Boiling Point Boiling point is the temperature at which the

vapor pressure of a liquid is equal to the external atmospheric pressure.

Remember on the last page that the vapor pressure of water at 100oC was 101.3 kPa or 1 atm, which is the external atmospheric pressure at sea level.

Therefore 100oC is the boiling point of water.

The boiling point varies with altitude. Why?

Because the external atmospheric pressure at altitude is lower (there’s less atmosphere at the top of the mountain!)

Normal boiling pt. = boiling pt. at 101.3 kPa

Boiling point

How much lower is the boiling point at high elevation?

In Denver, which is a mile high (1600m above sea level), the atmospheric pressure is lower than the 101.3 kPa at sea level, it is 85.3 kPa. In Denver, the boiling point of water is 95oC.

Boiling is also a cooling process, like evaporation is. It takes heat to make a water molecule convert from liquid to gas, and that heat leaves the system with the molecule.

What’s the difference between evaporation and boiling? Evaporation – Only involves molecules at

the surface that have enough kinetic energy to overcome (polar) attractions in the liquid to become gas.

Boiling – molecules in entire solution have enough kinetic energy to go from liquid to gas. Bubbles form all throughout the liquid and rise to the top as gas to escape the liquid.

Boiling Point Looking at this graph, state the boiling

point for each substance.

Section 10.3 The Nature of Solids Most solids are crystalline, which means they are

very ordered in a regular pattern and arranged by unit cells (smallest group of particles that maintains the pattern).

Solids In solids, particles are not as free to move.

The particles vibrate about a fixed point.

The particles do not flow or take the shape of the container.

The melting point – The temperature when a solid becomes a liquid (when the organization of its particles breaks down)

Ionic compounds have a higher melting point because they are crystals – they are orderly and have a 3D repeating pattern.

Solids melt when the vibrations become greater (due to thermal energy being added) than the forces holding the particles together.

Solids Some solids are not crystalline

If that is the case, they are amorphous.

Some solids may have the same formula, but they are connected differently = allotropes.

Diamond, graphite and buckyballs are C allotropes.

Amorphous substances

SiO2 , silicon dioxide, is anexample.

If hot molten SiO2 is rapidly cooled, it forms glass, which is amorphous, and sometimes called a supercooled liquid

Other forms of SiO2, when cooled slowly, form a crystalline structure. Quartz is one form of crystalline SiO2.

Some substances can be either amorphous or crystalline.

Section 10.4 Changes of State

Phase Diagrams

Solid Gas

Liquid

meltingfre

ezing condensation

vaporization

deposition

sublimation

Can you think of an example of sublimation? Dry ice Solid air fresheners Mothballs

Phase Diagrams A phase diagram gives the conditions of

temperature and pressure at which a substance exists as a solid, a liquid and a gas.

The blue area is solid. The green area is liquid.

The yellow area is gas. The lines show where

any two phases are in equilibrium (ice/water or water/steam)

The “Triple Point” at the bottom shows the only point where all 3 phases exist at one time.

Some interesting features

Check out how when the pressure is higher than atmospheric pressure (higher than 101.3 kPa), the melting point of water goes less than 0o C.

Ice skaters’ blades increase the pressure on the ice, causing the melting point to go lower than 0o C. That means a layer of liquid water forms under the skates, allowing the skaters to glide over the ice, rather than the skates scraping solid against solid.

Which line represents melting or freezing?

Which line represents sublimation or deposition?

Which line represents evaporation or condensation?

Quick Quiz

This is a normal phase diagram. Note how the slope of line B is positive. This is the normal situation for a substance where the density of the solid is more than the density of the liquid.

This is an unusual phase diagram (it is for water). That is because line B slopes backwards (has a neg. slope) because the density of ice is LESS than the density of water, so increasing pressure will make it melt, even if T stays the same.

Phase changes

As you heat up ice, and it converts to water and then to steam, during the phase changes, ALL added heat goes into creating the phase change, and the temperature doesn’t change.

Those are the two flat lines – the phase changes.

In the part labeled A, the H2O is in the form of ice. In the part labeled B, the ice is melting. In part C, all heat going into the system is used to

heat up the liquid water from 0 to 100oC. In part D, the water begins to boil, and the water

molecules start going into the gas phase (T constant). At E, the last molecule of water turns to steam, and

then all heat goes into heating up the steam >100oC.