The Kinetic Theory explains the effects of temperature and pressure on matter.

1. All matter is composed of small particles.
    
2. The particles of matter are in constant motion.
    
3. All collisions between the particles of matter are perfectly elastic.

   Elastic collisions transfer NO energy between particles. Each particle has exactly the same energy after the collision as before.

• An O₂ molecule in the air travels about 443 m/s (about 1700 km/hour) at 25 ^oC.
   
• The diameter of an O₂ is about 0.339 nanometers. Each molecule travels about 314 times its own diameter between collisions. At this rate, each oxygen molecule has over four and a half billion collisions per second.

• The SI unit for pressure is the pascal, named for Blaise Pascal.
   
• The pascal is a derived SI unit equal to 1 Newton / m²

  (1 N = (1 kg)(1 m) / s²)

• The average air pressure at sea level is 101.325 kilopascals (kPa).

  Scientists agree to call this average
  "standard atmospheric pressure".

• Kinetic energy is the energy an object has because of its motion.
   
• The SI units for energy is the joule.
   
• The joule is a derived SI unit equal to (kilograms) (meters) / s².
   
• KE = ¹/₂ mv²
   
• The calorie is sometimes used to describe heat. This unit is defined as the amount of heat needed to raise the temperature of one gram of water one Celsius degree.

  A food Calorie (notice the capital C) is 1000 regular calories, one Kilocalorie.

Measuring Temperature:

• The melting/freezing point of water at 1 atmosphere of pressure is 0 ^oC and 273 K.
   
• The boiling point of water at 1 atmosphere of pressure is 100 ^oC and 373 K.
   
• The interval between the melting/freexing point and the boiling point of water is divided into 100 intervals, each equal to 1 C^o and 1 Kelvin unit.
   
• The Kelvin scale extends downward from Celsius zero toward absolute zero - the lowest possible temperature. Therefore, the Kelvin scale has NO negative numbers.
   
• Kelvin temperatures are required when working "Gas Law" problems.
   
  • K = C^o + 273
     
  • C^o = K − 273

• A solid has a definite volume and shape.
• A liquid has a definite volume but not a definite shape.
• A gas has neither definite volume nor shape.

Use the Kinetic Theory to define the physical states of matter:

Solid - a substance whose particles have a low kinetic energy. The particles of a solid are held close together by intermolecular forces of attraction. Because the particles are so close together, they appear to vibrate around a fixed point.

When the temperature of a solid is raised, the velocity of the particles increases. The collisions between the particles occur with greater force, causing the particles to more farther apart. The ordered arrangement of the solid breaks down and a change in physical state occurs.

Liquid - a substance whose particles have enough kinetic energy to stretch the intermolecular forces of attraction. Collisions between the particles a strong enough to force the particles apart. The particles appear to have a moving vibration because they are still under the influence of the intermolecular forces of attraction.

As the temperature of a liquid is raised, the velocity of the particles increases. The collisions eventually become so great that the particles break all intermolecular forces, begin moving independently between collisions, and a change in physical state occurs.

Gas - a substance whose particles have enough kinetic energy to break all intermolecular forces of attraction. The particles of a gas move independently of each other. The particles move at random because they have overcome the intermolecular forces of attraction.

When a gas is raised to extreme temperatures, over 5000 ^oC, they have so much kinetic energy that their collisions will break electrons out of the atoms, and a change in physical state occurs.

Plasma - a charged gas. The particle collisions are violent enough to break electrons out of the atoms, producing particles with charges (electrons and positive ions).

Because of its extreme temperature, plasma is not common on Earth. Wet-lab chemistry is not concerned with plasma and its characteristics.

Physical state at room temperature (25 ^oC) and standard atmospheric pressure:

Under these conditions, the physical state of a substance is determined mainly by its chemical bond characteristics.

• Ionic compounds have strong electric charges holding the ions together as solids.
   
• Nonpolar molecular compounds of low molecular mass tend to be gases.
   
  • Greater molecular mass and greater polarity both tend to make substances more dense, producing either liquid or solid.

Within an atom, these forces are called weak forces, because they are much weaker than chemical bonds between atoms. Weak forces involve the attraction of the electrons of one atom for the protons of another atom.

When these forces interact between molecules, they are known as van der Waals forces, named for Johannes Diderik van der Waals.

The Kinetic Theory
 

Phase Changes

Every phase change is accompanied by a change in the energy of the system.

• As the temperature of a solid increases, the units of the solid vibrate with increasingly energetic motion. When the solid melts, the units that made up the solid are freed to move with respect to one another.

  This melting process is called fusion (different from nuclear fusion). The increased freedom of motion of the molecules or ions comes at a price, measured by the heat of fusion, or enthalpy of fusion, denoted ΔH_fus.
   

• As the temperature of a liquid increases, the molecules move with increasing energy. With increasing temperature, the concentration of gas-phase molecules just above the liquid increases. These gas-phase molecules exert a pressure on the liquid called vapor pressure.

  Vapor pressure increases with temperature until it equals atmospheric pressure. At this point the liquid boils. The energy required to cause this transition is called heat of vaporization or enthalpy of vaporization, denoted ΔH_vap.
   

• The molecules of a solid can be transformed directly into the gaseous state. The enthalpy change required for this transition is called the heat of sublimation, denoted ΔH_sub.

The graph above called a heating curve, representing the changes that occur when 1.00 mole of water is heated from −25 ^oC to 125 ^oC at a constant pressure of 1 atm.

Blue lines show the heating of one phase from a lower temperature to a higher one.

Red lines show the conversion of one phase to another at constant temperature.

Sample Problem

Calculating ΔH for Temperature and Phase Change:

What is the enthalpy change during a process in which 25.0 g (1.39 mol) of ice at −30 ^oC is converted to water vapor at 120 ^oC?

1. Raise temperature of ice from −30 ^oC to 0 ^oC.

   ΔH = m X s X &DeltaT = (25.0 g) (2.09 J/g-K) (30 K) = 1.56 kJ

2. Change phase from ice to water (liquid).

   ΔH = mol X ΔH_fus = (1.39 mol) (6.01 kJ/mol) = 8.35 kJ

3. Raise temperature of water (liquid) from 0 ^oC to 100 ^oC.

   ΔH = m X s X &DeltaT = (25.0 g) (4.18 J/g-K) (100 K) = 10.5 kJ

4. Change phase from water (liquid) to water vapor.

   ΔH = mol X ΔH_vap = (1.39 mol) (40.67 kJ/mol) = 65.5 kJ

5. Raise temperature of water vapor from 100 ^oC to 120 ^oC.

   ΔH = m X s X &DeltaT = (25.0 g) (1.84 J/g-K) (20 K) = 0.920 kJ

6. Total the enthalpy changes.

   ΔH_total = 1.56 kJ + 8.35 kJ + 10.5 kJ + 65.5 kJ + 0.920 kJ = 86.86 kJ

Heat Calculations

 
Vapor Pressure and Boiling Point

A liquid boils when the vapor pressure equals the external pressure acting on the surface of the liquid − in other words, when vapor pressure equals atmospheric pressure. At this point bubbles of vapor are able to form within the liquid.

The temperature at which a given liquid boils increases with increasing external pressure.

The boiling point of a liquid at 1 atm pressure is called its normal boiling point.
 

Determining Boiling Point

 
Phase Diagrams

A phase diagram is a graphical way to summarize the conditions under which equilibria exist between the different states of matter.

Phase diagrams are based on Gibbs' Phase Rule, stated by Josiah Willard Gibbs in the 1870's, and can be used to predict the stable phase of a substance at any temperature and pressure.

• Line AB is the liquid-vapor line, showing the vapor pressure of the liquid. It represents the equilibrium between the liquid and gas phase.
   
• Line AC is the solid-vapor line, representing the variation in the vapor pressure of the solid as it sublimes at different temperatures.
   
• Line AD is the solid-liquid line, representing the change in melting point of the solid with increasing pressure. This line usually slopes slightly to the right as pressure increases, because the solid phase of a substance is usually more dense than the liquid phase.
   
• On a phase diagram, a dotted line is drawn across the graph representing standard atmospheric pressure. The generalized graph here does not have that line.

• The point where the dotted line representing 1 atm crosses the liquid-vapor line is the normal boiling point of the substance.
   
• The liquid-vapor line ends at the critical point (B), which is the critical temperature and critical pressure of the substance. Beyond the critical point, the liquid and gas phases become indistinguishable from one another.
   
• The melting point of a substance is identical to its freezing point. The two differ only in the direction from which the phase change is approached. The melting point at 1 atm is the normal melting point.
   
• Where the three lines intersect (A), is known as the triple point. All three phases are in equilibrium at this temperature and pressure.

• Triple point (A), 0.0098 ^oC, 4.58 torr
• Normal melting point (B), 0 ^oC, 1 atm
• Normal boiling point (C), 100 ^oC, 1 atm
• Critical point (D), 374.4 ^oC, 217.7 atm

• Triple point (X), −56.4 ^oC, 5.11 atm
• Normal sublimation point (Y), −78.5 ^oC, 1 atm
• Critical point (Z), 31.1 ^oC, 73.0 atm
• Notice that CO₂ does not exist in the liquid phase at 1 atm of pressure.

Computer Lab