Chapter 13

States of Matter

I. Gases are one of the four states of matter. Solid, Liquid, Gas, and Plasma

A.  Gases share 5 characteristics

  1. Expansion – indefinitely and uniformly to fill all the space in a container.
  2. Indefinite shape or volume takes the shape of the container.
  3. Compressibility gases can be highly compressed.
  4. Low density – much lower than solids or liquids
  5. When mixed two non-reacting mix completely and uniformed.

B.  When applied to gasses, this theory is called Kinetic Molecular Theory

  1. Gases are composed of molecules and the distance between these molecules is large compared to the size of the molecules
  2. No attractive forces exist between the molecules in a gas. This is what keeps a gas from becoming a liquid.
  3. The molecules are in a constant state of rapid motion.
  4. All collisions are perfectly elastic. This means there is no loss of kinetic energy. Kinetic energy is the energy of motion.
  5. The average kinetic energy per molecule is proportional to the temperature in Kelvin. The average kinetic energy per molecule is the same at a given temperature for all gases regardless of the mass of the gas. Theoretically, at absolute zero all motion stops.

O Kelvin = -273 oC

C.  Gases conform to the assumption called Ideal Gases

  1. Ideal Gas Law assumptions
  2. The molecules in ideal gas are considered to be dimensionless points that have no volume.
  3. The molecules of a gas are in constant straight-line motion.
  4. The collisions are perfectly elastic.
  5. The molecules of an ideal gas do not exert any attractive forces on each other.
  6. Real gases behave as ideal gases under moderate conditions, but deviate if the temperature is very low or the pressure is very high.

II.  The Liquid State

A.  Liquids are one of the four states of matter.

There are 6 characteristics:

1.  Limited expansion.

2.  Shape – they have no characteristic shape; they take the shape of the container.

3.  Volume – they maintain their volume.

4.  Compressibility – slightly compressible.

5.  High density – when compared to gases.

6.  Mixing is slow compared to gases.

B.  Unlike gases molecules in liquids are close together and do exert some attractive or repulsive forces upon one another and the collisions are not perfectly elastic.

III.  Condensation and evaporation

A.  Real gases can be liquefied when attractive forces overcome the kinetic energy of the molecules. This is achieved two ways:

1.  Decreasing the temperature

2.  Increasing the pressure, significantly.

B.  Condensation is the conversion of a gas to a liquid.

C.  Evaporation is the conversion from a liquid to a gas.

1.  Opposite of condensation.

IV.  Vapor Pressure

A.  If you fill a closed container half-way with a liquid, some of the liquid will evaporate and be trapped in the space between the liquid and the lid.

1.  This will reach maximum allowed point for a specific temperature.

2.  In addition, some of the gaseous liquid will return to the liquid state and more liquid will go into the gaseous state.

3.  When these two processes (liquid à gas and gasà liquid) are equal, we have a dynamic equilibrium.

B.  When this dynamic equilibrium is established, the concentration of molecules in the gas form remains constant at constant temperature and the vapor exerts a definite constant pressure called vapor pressure of the liquid. This is at a fixed temperature.

C.  Example: compare gasoline and water. Gas has a higher vapor pressure than water.

v  Gas would evaporate faster than water.

And at the same temperature, gas would have more molecules in the gas form than water.

V.  Boiling point

A.  When the vapor pressure equals the external pressure above the surface of the liquid, bubbles of vapor forms rapidly throughout the liquid and the liquid boils.

B.  Since atmospheric pressure varies with altitude and atmospheric conditions, scientists report boiling points at a standard pressure of 760 mm Hg or 1 atm.

C.  Evaporation and boiling point:

1.  Evaporation of a liquid at its boiling point involves a loss of energy by the liquid. So, heat must be supplied for the temperature to remain constant.

2.  It is an endothermic change of state requiring the absorption of energy.

v  Condensation is an exothermic change of state requiring a release of heat.

This is why steam burns you more than water. Steam condenses on the cool skin releasing heat, which cooks the tissue and burns you more.

VI.  Distillation

A.  Distillation is the process of purifying of a liquid by heating it to the boiling point and cooling the vapors in a condenser. This process takes advantage of the different boiling points of different liquids.

B.  Common applications:

1.  Separating water from dissolved salts – the purified compound is called the distillate and the impure stuff remaining is called the residue.

Fractional distillation – many simple distillations. Petroleum industry.

VII.  Surface tension and viscosity.

A.  Surface tension – the property of a liquid that tends to draw the surface molecules into the body of the liquid and reduce the surface to a minimum.

1.  As temperature increases, surface tension decreases.

2.  The reason why we read the bottom of the meniscus on a graduated cylinder is that water molecules pulls itself down into the body of the water.

B.  Viscosity – the property of the liquid that describes the resistance of a liquid flow.

1.  Highly viscous – honey

2.  Medium viscous – water

3.  Low viscosity – gas

VIII.  The solid state

A.  Solids have particles that are closer together than liquids, except water.

B.  They are subject to strong attractive and repulsive forces between the molecules.

C.  Particles have a relatively fixed position with respect to one another.

D.  6 general characteristics:

1.  No expansion at a constant temperature.

2.  Definite shape.

3.  Maintain their volume.

4.  Practically incompressible.

5.  High density – when compared to gases.

6.  Severely limited mixing.

IX.  The shape of solids.

A.  Crystalline – any solid that consists of particles arranged in a definite geometric shape or form that is distinctive for that solid.

1. Example – diamond or quartz.

B.  Amorphous – particles arranged in an irregular manner and lacks the order found in crystals.

1.  Example – glass and plastics.

X.  Melting or freezing point

A.  Freezing point – the temperature at which the particles of a liquid begin to form crystals or irregular particles of a solid. When something freezes this is an exothermic process.

B.  Melting point – the temperature at which a kinetic energy of some of the particles in a solid matches the attractive forces in the solid and the solid begins to liquefy. When something melts this is an endothermic process.

XI.  Sublimation

A.  Sublimation – direct conversion of a solid to a gas without passing through the liquid state. Þ endothermic.

B.  Deposition – the direct conversion of gas to a solid without passing through a liquid state. Þ exothermic.

Melting Vaporization

SOLID LIQUID GAS

Freezing Condensation

Sublimation

SOLID GAS

Deposition

If the arrow goes to the left it is exothermic. If the arrow goes to the right it is an endothermic process.

XII.  Heat energy transformations in the three physical states of matter.

A.  As you heat up a solid, the kinetic energy of the particles of the solid increases until you get to the melting point.

B.  At the melting point of the solid, you need to continue to add more heat to melt the/more solid. This is because it is an endothermic process.

C.  You get a liquid. As you heat the liquid, the kinetic energy increases until you reach the boiling point of the liquid.

D.  At the boiling point of the liquid, you need to continue to add more heat to boil more liquid. This is because it is an endothermic process.

E.  You get a gas. Further heating of the gas increases the kinetic energy of the molecules and the temperature increases.

v  Example – Calculate the quantity of heat in calories required to convert 120. g of ice at 0oC to steam at a 100oC.

v  We will be going from (a) a solid to a liquid (heat of fusion), (b) a cold liquid to a hot liquid (specific heat), and (c) and from a liquid to a gas (heat of vaporization)

a)  Hfusion water = 80 cal/g.

b)  Specific heat of water = 1 cal/g oC

c)  Hvaporization water = 540 cal/g.

a) 120 g x 80 cal = 9600 cal

g

b)  1 cal x 120 g x (100 – 0) oC = 12000 cal

g oC

c) 120 g x 540 cal = 64800 cal

g

d) Add

(9600 + 12000 + 64800) cal = 86400 cal.

XIII. Intermolecular Forces and Intramolecular Forces

A.  Intermolecular Forces are forces of attraction between neighboring molecules.

1. Ionic bonds result in the formation of a positive and a negative charge on the molecule by transferring their valance electrons. These opposite charges are a source of attraction and result in the formation of weak bonds between the molecules.

2. Hydrogen bonding this is a bond that forms when there is a Hydrogen and either a F,O, or N. This bond is about 15% as strong as the covalent or ionic that holds the molecule together.

3. Metallic bonds atoms of metal achieve a more stable arrangement by sharing their valance electrons. This type of bonding is the reason that most metals are solids at room temperature.

B.  Intramolecular Forces are forces that exist inside the molecule. These forces are usually very weak, but when added together these are the forces that hold DNA together.