KINETIC-MOLECULAR THEORY



1. SC6. Students will understand the effects motion of atoms and molecules in chemical and physical processes.

1. a. Compare and contrast atomic/molecular motion in solids, liquids, gases, and plasmas.

2. b. Collect data and calculate the amount of heat given off or taken in by chemical or physical processes.

3. c. Analyzing (both conceptually and quantitatively) flow of energy during change of state (phase).

KINETIC-MOLECULAR THEORY

The Kinetic-Molecular theory explains the effects of temperature and pressure on matter. It is described in terms of particles in motion.

1. All matter is composed of small particles (atoms, ions, molecules). The phase of matter determines the space between the particles.

2. The small particles are in constant motion. They move in a straight line until they collide with another particle or the wall of their container. All collisions between these particles are perfectly elastic – meaning that there is no change in the total kinetic energy of the two particles before and after the collision.

3. The energy of the particle can be determined by mass and velocity. KE = ½mv2 where m=mass and v=velocity. All particles do not have the same kinetic energy.

In the demonstration with the cold and hot water, which temperature water mixed the dye the fastest? Why did it? Explain.

TEMPERATURE AND HEAT

Average kinetic energy is the average of the kinetic energies of all the particles of matter in a given sample. It would be an impossible task to measure the kinetic energy of each particle directly (you are breathing in millions of molecules of oxygen right now – could you measure the energy of EACH one?!). So, temperature is used.

The TEMPERATURE of a substance is a measure of the average kinetic energy of the particles in a sample of matter. Since kinetic energy is conserved, the average kinetic energy of the particles remains the same and can be measured. Therefore, if the temperature goes down, then the average kinetic energy of the particles has also gone down and they move more slowly.

Temperature can also be used to determine the flow of energy. Warmer objects will have a higher kinetic energy than cooler objects. Energy is ALWAYS transferred from a warmer object to a cooler object. The amount of energy transferred is called HEAT.

PHASES OF MATTER

The phase of matter (solid, liquid, gas, plasma) is determined by two factors:

1. Temperature (kinetic energy) of particles.

2. Intermolecular Forces (IMF) of Attraction - These are the forces between particles.

a. Dispersion forces – between nonpolar molecules

b. Dipole-Dipole forces – between polar molecules

c. Hydrogen bonding – between special polar molecules

All INTERmolecular forces are weaker than INTRAmolecular foces which are true bonds (ionic, covalent, metallic).

CHARACTERISTICS OF THE PHASES OF MATTER

Solids: definite shape, definite volume, strong IMF, particles are in a fixed position and vibrate around a fixed point.

Liquids: no definite shape, definite volume, strong IMF, particles can flow past each other but cannot break away completely.

Gas: no definite shape, no definite volume, very weak IMF, particles are independent of each other and can go wherever they want inside their container. Another word for gas is vapor.

Plasma: no definite shape, no definite volume, composed of electrons and positive ions. The electrons have been knocked away from the atoms due to the high temperatures.

Looking at the boxes below and pretend that they are beakers sitting on a counter. Show what the atoms of a solid, liquid, and gas would look like in each.

solid liquid gas

INTERMOLECULAR FORCES (IMFs)

Dispersion Forces – a weak force that results in the attraction of normally non-polar particles; the weakest intermolecular force; when nonpolar molecules collide with each other, the electron clouds are temporarily shifted, creating a temporary dipole; The more electrons in the electron cloud, the stronger the dispersion force. This is why iodine is a solid, bromine a liquid, and chlorine and fluorine are gases at room temperature.

1. Dipole-dipole forces – attractions between oppositely charged regions of polar molecules; polar molecules have permanent dipoles due to the fact that some regions are always positive and some regions are always negative; these forces are stronger than dispersion forces if looking at particles of the same mass.

2. Hydrogen bonds – (not a bond); a special dipole-dipole attraction; occurs when the molecule contains hydrogen and the hydrogen is covalently bonded to a highly electronegative atom (only F, O, N); the electronegative atom has almost complete possession of the electron pair shared with the hydrogen. The hydrogen has a strong partial positive charge. The molecule is highly polar; the O-H bonds in water are more polar than the N-H bonds in ammonia – therefore the hydrogen bonds between water molecules are stronger than those between ammonia molecules. Both have higher than expected boiling points (strong IMF).

MORE ON LIQUIDS

1. Density– denser than gases

2. Compression – enormous pressure can make liquids compress a small amount.

3. Fluidity – ability to flow; gases and liquids are fluids because they can flow through each other. Liquids are slower to diffuse than gases.

4. Viscosity – a measure of the resistance of a liquid to flow; determined by the intermolecular forces, shape of particles, and temperature.

5. Surface tension – it is a measure of the inward pull by particles in the interior. The stronger the particle’s attraction for each other, the stronger the surface tension; Water has a high surface tension due to the attractive forces of the water molecules; Surfactants reduce surface tension.

6. Capillary Action – the adhesion or attraction between molecules that are different; can cause the meniscus in the graduated cylinder or the wicking property of paper towels. Cohesion = attraction to each other. Adhesion = attraction to another.

A SPECIAL CASE: WATER

Water’s unique properties are a result of hydrogen bonding.

Phases of Water:

Solid: Each hydrogen is bonded to four other water molecules in a open crystalline shape.

Liquid: The solid crystal lattice collapses; the liquid molecules flow closer to each other and the liquid is more dense than ice as a result. Ice floats!

Surface Tension:

Particles in the interior of the liquid are subjected to attractive forces in all directions. Molecules at the surface have a new inward attraction that results in surface tension. Liquids form spheres when dropped because spheres have the least amount of surface area.

Capillary Rise (action):

This is the rise of a liquid in a tube of small diameter. An attractive force between the tube wall and the liquid will cause the liquid to rise.

MORE ON SOLIDS

1. Density – more closely packed than a liquid, thus they are denser; usually a 10% difference between the solid and liquid density of the same substance; Water is the exception.

2. Compression – enormous pressure can make solids compress a small amount.

3. Types of solids:

a. Crystalline Solids – a solid whose particles are arranged in an orderly, geometric, three-dimensional structure; examples many minerals.

b. Molecular Solids – covalent compounds generally have weak attractions, so large molar masses are required in order to be a solid; examples include the sugars.

c. Covalent Network Solids – made from atoms that form multiple covalent bonds; examples are carbon and silicon.

d. Ionic Solids – have network of attractions between the negative and positive ions that allow for a greater attraction throughout the crystal.

e. Metallic solids – a solid formed from positive metal ions surrounded by a sea of electrons; strength varies, but is usually more than covalent compounds.

f. Amorphous solids – a solid where the particles are not arranged in a crystalline pattern; Often forms when a molten (melted solid) cools too quickly to allow time for the crystals to grow; Examples include glass, rubber, and many plastics.

PHASE CHANGES

Heat is energy that causes the particles of matter to move faster and further apart. Therefore, the particles can then change phases of matter. Adding heat increases the temperature.

Phase changes are accompanied by a change in heat energy, but not temperature. Heat energy is used to overcome forces that hold the particles together. Phase changes produce changes in physical properties only.

Solid ( Liquid melting

Liquid ( Gas vaporization endothermic

Solid ( Gas sublimation

Gas ( Liquid condensation

Liquid ( Solid freezing exothermic

Gas ( Solid deposition

Substances are made to change phase by adding or taking away energy. When undergoing a phase change, the mass remains constant and volume changes. Thus, density changes.

VAPOR PRESSURE

Vapor pressure is the pressure exerted by a vapor (gas) in equilibrium with its liquid. Intermolecular forces (IMF) determine vapor pressure. (Remember IMF - the forces of attraction between neighboring molecules). The lower the vapor pressure, the stronger the IMF. Doesn’t that make sense? The stronger the IMF, the harder it is for the liquid to become a gas, therefore the vapor pressure is lower.

What is a vapor? How does a vapor differ from a gas?

VAPORIZATION AND CONDENSATION

Vaporization is the change in state from a liquid to a gas - includes evaporation and boiling. Evaporation occurs when a liquid molecule gets enough kinetic energy to leave the surface of a non-boiling liquid. It gets enough energy to overcome the IMF within the liquid. It’s like leaving a glass of liquid water out on a counter top and it eventually all turns into a gas.

Condensation is the change in state from a gas to a liquid. The IMF traps any particle colliding with the surface of a liquid.

How are boiling and evaporation alike? How are boiling and evaporation different?

Liquid-Vapor Equilibrium

In a closed container half-filled with a liquid, the liquid will evaporate into the space above the liquid. Soon, molecules will also condense back to the surface of the liquid.

Rate of vaporization

R

a Rate of condensation

t

e

In a closed system, a liquid and its vapor will reach an equilibrium at a specific pressure for a particular temperature (i. e. all particles have the same average kinetic energy). The rate of vaporization is equal to the rate of condensation. A state of dynamic equilibrium is reached. Change is still going on, but the overall effect does not change.

Why is it called dynamic equilibrium?

Equilibrium Vapor Pressure

The molecules in dynamic equilibrium exert a pressure called equilibrium vapor pressure - the pressure exerted by a vapor in equilibrium with its liquid. Vapor pressure will increase steadily as temperature increases. This indicates there are a greater number of molecules present as a vapor. Initial vapor pressure is determined by the intermolecular forces. Examples of substances’ initial vapor pressure are: Mercury - 0.0002 kPa, Water - 3.167 kPa, Acetone - 30.8 kPa.

Substances with low vapor pressures have strong IMF. (Ionic compounds have very low vapor pressures.) Substances with high vapor pressures have weak IMF. Mercury is a very dense liquid with high IMF. It has a very low initial vapor pressure. Acetone is a volatile liquid – it changes into a vapor very easily and has low IMF. Therefore, its initial vapor pressure is high (more molecules are in the vapor state).

Boiling Point

As temperature increases, the vapor pressure of the liquid increases because the kinetic energy of the molecules increases. When KE increases enough to overcome the internal pressure of the liquid caused by the pressure of the atmosphere on the liquid’s surface, the molecules collide violently enough to push each other apart. When the vapor pressure is equal to atmospheric pressure, the liquid boils.

Bubbles form due to this pushing apart and they rise to the surface. (They are less dense than the liquid) and the liquid boils. Boiling takes place throughout the liquid.

The NORMAL BOILING POINT is the temperature at which the vapor pressure is equal to one atmosphere (101.3 kPa). Boiling point is a function of pressure. The lower the pressure is, the lower the boiling point is.

Define boiling point in terms of vapor pressure:

MELTING AND FREEZING

Freezing and melting require less change in energy than vaporization and condensation. The atoms/molecules are already close together. The Freezing point is the temperature at which the vapor pressure of the solid and the vapor pressure of the liquid are equal. (Also called the melting point.) It is not affected much by a change in external pressure. But it is dependent upon the IMF of the substance. A weak IMF means a low melting point.

What happens as an ice cube sits on a table in a 22(C lab? Is equilibrium established? Explain.

SUBLIMATION AND DEPOSITION

Solids with a high vapor pressure (at room temperature) go straight from solid to gas, bypassing the liquid phase. This is called sublimation. The opposite change from a gas to a solid is called deposition.

Give three examples of substances that sublime readily.

SUMMARY

Strong IMF Weak IMF

Nonvolatile volatile

low evaporation rates high evaporation rates

high boiling point low boiling point

low vapor pressure at room temperature. high vapor pressures at room temperature.

PHASE CHANGES

Heat is energy that causes the particles of matter to move faster and further apart. Therefore, particles can then change phases of matter. Adding heat increases temperature. Phase changes are accompanied by an increase in heat energy, but not temperature.

Heat energy is used to overcome the forces that hold the particles together. Phase changes produce changes in physical properties only.

A HEATING CURVE shows how temperature and energy are related. A PHASE DIAGRAM shows how temperature and pressure are related.

How are heating curves and phase diagrams alike? How are they different?

HEATING CURVE OF WATER

E

D

C

B

A

TIME

A. As heat is supplied, the heat increases the KE of the molecules.

B. When the substance reaches the melting point, all heat being supplied is used to change the phase (potential energy change). Temperature will not change until the phase is complete.

C. When the phase change is complete, heat again raises the KE of the molecules, temperature changes and increases to the boiling point.

D. All heat is used the change the phase. Same as “B”.

E.

PHASE DIAGRAM FOR WATER

Tm - normal melting point at STP

Tb - normal boiling point at STP

Triple Point - all three phases are at equilibrium

D - critical temperature – the highest temperature for a vapor and liquid at equilibrium;

Also critical pressure - pressure needed to liquefy gas at the critical temperature.

Solid lines are the equilibrium lines between phases.

[pic]

ENERGY

Energy is the ability to do work or produce heat. There is potential energy and kinetic energy. Potential energy is stored energy and kinetic energy is the energy of motion.

Remember that energy is conserved – it can be converted from one form to another, but it cannot be destroyed.

The energy stored in a substance because of its composition is CHEMICAL POTENTIAL ENERGY. Chemical potential energy plays an important role in chemical reactions. In a chemical reaction, the potential energy can be released as heat. HEAT is the energy that flows from a warmer object to a colder object.

Heat is measured using the SI unit JOULE, J. (The other unit is calorie. One calorie = 4.184 joules.)

ENERGY AND CHANGES OF STATE

When heat is added to a solid, the KE of the molecules increases. When the melting point is reached, the added energy (heat) increases the KE of the molecules. The positions of the molecules are changed. An equation to determine the amount of heat is the following:

q = (m) (∆T) (Cp)

where “q” is heat, “m” is mass,

“∆T” is change in temperature, and

“Cp” is the specific heat of the substance you are measuring.

Specific heat of any substance is the amount of heat required to raise the temperature of one gram of that substance by one degree Celsius. Cpice - 2.06 J/g•C° Cpwater - 4.18 J/g•C° Cpvapor - 2.02 J/g•C°

Sample Problem

Calculate the amount of energy in joules needed to raise 100.0 g of ice from –75.0°C to 0.00°C.

q = (m)(∆T)(Cpice)

q = (100.0 g)(75.0°C)(2.06 J) = 15,450 joules

g • C° = 15,500 J

Practice Problem

Calculate the amount of energy in joules needed to raise 100.0 g of ice from 25.0°C to 75.0°C.

Enthalpy, H, of a substance is its energy plus a small added term that takes into account the temperature and pressure of the substance; the enthalpy and the energy of a substance are very close to each other; at a constant pressure, enthalpy is equal to the heat (energy) that is transferred between reactants and products

Enthalpy (Heat) of Fusion, ∆Hfus

Energy required to melt one gram of a specific substance at its melting point; the energy of melting.

Enthalpy (Heat) of Vaporization, ∆Hvap

Energy required to vaporize one gram of a substance at its boiling point; the energy of boiling; directly related to the strength of the intermolecular forces that exist in the liquid. High IMF means low vapor pressure and high ∆Hvap . (This is because much energy is needed to increase the kinetic motion of individual molecules to free them from intermolecular attraction.)

Constants:

∆Hfus of water - 334 J/g

∆Hvap of water - 2260 J/g

Sample Problem:

How much heat is required to melt 5.67 g of iron (II) oxide, FeO, if its enthalpy of fusion is 450.0 J/g?

heat mass

q = m • ∆Hfus

q = m • ∆Hfus

q = (5.67 g) (450.0 J) = 2551.5 = 2.55 x 103 J

(g)

Sample Problem:

How much heat is required to melt 50.0g of water at its melting point?

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