Gases and States of Matter: Unit 8

States of Matter

There are three states (also called phases) of matter.

The picture represents the same chemical substance, just in different states.

There are three states (also called phases) of matter.

• Solid: Matter that has both a definite shape and definite volume. Molecules or atoms are very close together and can only vibrate a little. They do not move past each other. Solids are not easily compressed and have a rigid shape and structure.

There are three states (also called phases) of matter.

• Liquid: Matter that has a distinct volume but no specific shape. Molecules or atoms are close together but have the ability to slide across one another very easily. Liquids are more compressible than solids. Their shape will change as the container changes while their structure remains the same.

There are three states (also called phases) of matter.

• Gas: matter that has no fixed volume or shape. It conforms to the volume and shape of its container. Its molecules or atoms are very far apart from each other and move very fast. Gases are extremely compressible. Gases will also take the shape of their container while their the structure stays the same.

Density Comparison

• If you consider the solid, liquid, and gas state of one particular substance, this rule holds true in most cases:

• Solid is more dense than liquid

• Liquid is more dense than gas

Water is Weird…

• A notable exception is water! The solid state of H2O, ice, is less dense than liquid water. This is why ice floats.

• This is true because of the way hydrogen bonds form when liquid water freezes. The hexagonal pattern results in empty space between the molecules.

Two Types of Solids

• Crystalline Solids

• Molecules are packed together in a predictable way. They are arranged in an orderly, geometric, three dimensional structure.

• The smallest repeating part of a crystalline structure is called a unit cell. The measured strength of the bond that holds these crystals together is lattice energy.

• Ex: cubic, hexagonal, rhombohedral, etc.

Ionic solids

• Crystal lattice is formed from alternating anion and cation.

• High melting point and extremely hard.

• Always solids at RT.

Covalent network solid• Form a 3-D covalent network; very strong with

high melting point.

• Carbon and silicon

Two Types of Solids

• Amorphous Solids

• Particles are NOT arranged in a regular repeating manner.

• Amorphous means “without shape.”

• Examples: Glass, rubber, plastics, wax,

Molecular solids

• Units are molecules, held together by weak IMFs. Low melting points.

• Ice, dry ice and sugar (most are not solid at RT)

Atomic solids

• Unit particles are atoms

• Noble gases when they are cooled to solid state. Usually very soft because they have weak IMFs.

Metallic solids

• Atoms are surrounded by mobile valence electrons.

• Metals are malleable, ductile and good conductors.

Liquids

• Fluidity – liquids (and gases) have the ability to flow.

• Viscosity – the measure of the resistance of a liquid to flow. Cold pancake syrup is very resistant to flow – it is viscous. Water flows easily – it is less viscous.

Liquids

Viscosity

• Viscosity decreases with increasing temperature. Ex: heating up the syrup in the microwave makes it pour easier.

Buoyancy

• Buoyancy is the upward force that keeps things afloat. This force enables the object to float or at least seem lighter.

Phase changes

• Matter can change from one phase to another by adding or removing energy.

• There are six phase changes.

Phase Changes That Require Energy (increase in energy, endothermic)

• Melting: solid liquid

• Ex: ice melting to liquid water

• Vaporization: liquid gas

• Ex: heating water on the stove, steam is released.

• Sublimation: solid directly to gas

• Ex: dry ice (solid CO2) lets off CO2 gas, looks like steam; solid air fresheners.

Phase Changes that Release Energy (decrease in energy, exothermic)

• Condensation: gas liquid

• Ex: “sweat” collects on a cold glass of ice water.

• Freezing: liquid solid

• Ex: water freezing into ice.

• Deposition: gas directly to solid

• Ex: frost forms on the grass on a cold morning, snowflakes form from water vapor in clouds.

Boiling vs. Evaporation

• Boiling – heating a liquid to the temperature at which all molecules have enough energy to escape and vaporize.

• Evaporation – vaporization of surface molecules, very slow. Does not occur at high temperatures.

Phase diagrams

• A phase diagram shows what phase a substance will be in at a certain temperature and pressure. Pressure is measure in atmospheres (atm).

Phase Diagram

• Triple Point – is the point on a phase diagram that shows the temperature and pressure combination at which three phases of a substance can coexist.

Phase Diagram

►Critical Point – the temperature and pressure combination above which a vapor cannot be liquefied under any circumstances.

Critical point

• Compare the phase diagram of water and CO2. The solid-liquid line for water has a negative slope. The solid-liquid line for CO2

has a positive slope. What does that mean?

• This information tells us that water’s solid phase is less dense than its liquid state. CO2’s solid phase is more dense than its liquid state. This is because increasing pressure favors the more dense phase.

1) At which pressure and temperature do all three phases coexist (triple point)?

2) Give a possible pressure and temperature combination for a solid.3) Is it possible for this substance to be liquid at -60 degrees Celsius?

Why or why not?

Heating curves• When energy/heat is added to or removed from a substance, the

following could happen:– Temperature changes (molecules change speed)

– Phase change

Q for heating/adding energy is always positive (endothermic)Q for cooling/releasing energy is always negative (exothermic)

How much energy is needed to convert 153 grams of ice at -15°C to steam at 125°C? The molar mass of water is 18.016 g/mol.