States of Matter

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# States of Matter - PowerPoint PPT Presentation

States of Matter. Mr. Solsman Chapter 12. What makes solids, liquids, and gases different?. What are the macroscopic differences? What are the molecular level interactions that cause these differences? What are the molecular level changes that must happen for phase change to happen?

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### States of Matter

Mr. Solsman

Chapter 12

What makes solids, liquids, and gases different?
• What are the macroscopic differences?
• What are the molecular level interactions that cause these differences?
• What are the molecular level changes that must happen for phase change to happen?
• What are the energy changes that accompany phase changes?
KINETIC THEORY OF GAS BEHAVIOR

1. All gases are composed of tiny particles called molecules.

2. These molecules are in constant random motion.

3. On an average, these molecules are far apart.

4. If the temperature cools, molecular velocities decrease. Inother words, the average kinetic energy goes down.

5.Molecules are perfectly elastic. They experience no net loss of energy on collision.

6.There are attractive forces between the molecules of a gas.

Explaining the Behavior of Gases
• The KMT helps explain the behavior of gases.
• Why do gases expand until they fill a container?
• Why are gases compressible and liquids not?
• How do odors make their way across a room (perfume)?

Gases easily flow past each other because there are no significant forces of attraction.

• Diffusionis the movement of one material through another.
• Effusion is a gas escaping through a tiny opening.

Graham’s law of effusion states that the rate of effusion for a gas is inversely proportional to the square root of its molar mass.

Ammonia has a molar mass of 17.0 g/mol; hydrogen chloride a molar mass of 36.5 g/mol. What is the ratio of their diffusion rates?

• Calculate the ratio of diffusion rates for carbon monoxide and carbon dioxide.
Gas Pressure
• Pressureis defined as force per unit area.
• Gas particles exert pressure when they collide with the walls of their container.
• P = F / A

The particles in the earth’s atmosphere exert pressure in all directions called air pressure.

• There is less air pressure at high altitudes because there are fewer particles present, since the force of gravity is less.

The SI unit of force is the Newton (N).

• One pascal(Pa) is equal to a force of one Newton per square meter or N/m2.
• One atmosphereis equal to 760 mm Hg or 101.3 kilopascals.

Dalton’s law of partial pressures states that the total pressure of a mixture of gases is equal to the sum of the pressures of all the gases of the mixture.

• The partial pressure of a gas depends on the number of moles, size of the container, and temperature and is independent of the type of gas.

General formula:

Pt = P1 + P2 + P3 + ……+ Pn

A mixture of oxygen, carbon dioxide, and nitrogen has a total pressure of 0.97 atm. What is the partial pressure of oxygen if the partial pressure of carbon dioxide is 0.70 atm and the partial pressure of nitrogen is 0.12 atm?

Suppose 10.1 g of Ne, 8.00g O2 , and 2.80 g N2 are mixed in a flask. If the total pressure is .500 atm, what is the partial pressure of each gas?

If 16.0 grams of oxygen, 2.0 grams of hydrogen, 8.00 grams methane, and 8.0 grams of He are mixed for a total pressure of 4.00 atm, what are their individual partial pressures?

What makes a Gas Condense?
• When molecules collide there is an attractive force that makes them want to stick together.
What makes a Gas Condense?
• If the kinetic energy is high and the force low, they separate and remain as a gas.
What makes a Gas Condense?
• Lower the temperature and the gas molecules slow down. Eventually the kinetic energy is not strong enough to overcome the attractive force.
• The molecules begin to stick together and form a liquid.
What makes a Gas Condense?
• If you cool a gas, eventually the molecules will begin to stick together.
• If the force is strong this happens at a relatively high temperature.
• If the force is weak this happens at a very low temperature.
From the other direction:
• Imagine that you have two cold liquids.
• Liquid A has strong attractive forces.
• Liquid B has weak attractive forces.
• As you heat these liquids, which one will vaporize at the lower temperature?
KINETIC THEORY OF A LIQUID

All liquids are composed of clusters of particles weakly bonded* to each other but free enough to move over one another**.

* This explains the definite volume.

**This explains no definite shape.

What Are The Forces?
• In Pure Liquids
• Dipole – Dipole
• London Dispersion Forces
• Aka Instantaneous dipole-induced dipole
Dipoles
• A dipole exists when the atom at one end of a chemical bonds has a stronger attraction for the electrons being shared.
• That creates a positive (δ+) end and a negative (δ-) end.
• Opposite charges attract each other.
• This is a dipole-dipole attraction.
Hydrogen “Bonds”
• A hydrogen “bond” is a very strong dipole-dipole attraction that forms when molecules have a δ+ on a hydrogen and an unshared pair of electrons on an atom with δ- . This second atom is always an oxygen, nitrogen, or fluorine atom.
Nonpolar Liquids
• Instantaneous dipole-induced dipole attractions are the weak attractions created by temporary dipoles in nonpolar molecules causing (inducing) a weak temporary dipole in a nearby molecule.
Boiling Temperature
• The stronger the forces between molecules, the higher we must heat them to overcome these forces.
BOILING TEMPERATURES OF THREE SUBSTANCES

FORMULA MASS_TC_

CH4 16.0 -161 0C

H2O 18.0 +100.0 0C

CO2 44.0 -56.6 0C

As attractive forces become greater, the amount of energy needed to overcome the force is greater and the boiling temperature is higher.

Kinetic vs Potential Energyin Vaporization
• If you have a liquid, the molecules are held together by some attractive force.
• To separate these molecules against that force requires energy.
• This process gets energy from kinetic energy and stores it as potential energy.
THE KINETIC THEORY OF SOLIDS

All solids are composed of one giant cluster or aggregate of particles strongly bonded to each other in fixed positions*, but free enough to vibrate and rotate about those positions.

* This accounts for the definite volume and shape of a solid.

Contrasting the molar volume of a gas to that of its solid.

A.The molar volume of N2(g) at STP is 22.4 liters/mole.

B.The molar volume of N2(s) at -2100C is 27.2ml/mole.

PHASE CHANGES

Phase Changes are physical changes. There is no loss in identity. The molecular, atomic, or empirical formulas do not change.

Na+Cl-(s) Na+Cl-(l) - MELTING

H2O(l)  H2O(g) -EVAPORATION

CO2(s)  CO2(g) -SUBLIMATION

Ne(g)  Ne(l) - CONDENSATION

H2O(g)  H2O(s) -DEPOSITION

Cu(l)  Cu(s) -FREEZING

MOLAR HEAT OF VAPORIZATION – the amount of heat needed to vaporize one mole of a liquid at its boiling point.

MOLAR HEAT OF MELTING(FUSION) – the amount of heat needed to melt one mole of a solid at its melting point.

TYPES OF ENERGY CHANGES
• Exothermic changes – changes where energy is released. For chemical and physical changes, energy is expressed as one of the products.
• Endothermic changes – changes where energy is absorbed. For chemical and physical changes, energy is expressed as one of the reactants.
ENDOTHERMIC PHASE CHANGES

1. Melting 1.44 kcal + H2O(s)  H2O(l)

2. Evaporation 4.88 kcal + Cl2(l)  Cl2(g)

3. Sublimation 18.1 kcal + CO2(s)  CO2(g)

In each of these: Low PE  High PE

As energy is absorbed from the surroundings at constant temperature, bonds break and potential energy rises. These produce a cooling effect.

EXOTHERMIC PHASE CHANGES

Freezing H2O(l)  H2O(s) + 1.44kcal

Condensation Ne(g)  Ne(l) + 0.405kcal

Deposition H2O(g)  H2O(s) +11.3kcal

In each of these: High PE  Low PE

Potential energy drops as bonds form at constant temperature. Energy is released to the surroundings. These produce warming effects.

CONTRASTING HEATS OF FUSION AND HEATS OF VAPORIZATION

Substance Heat of Fusion Heat of Vap

Ne 0.080kcal 0.405kcal

H2O 1.44kcal 9.7kcal

NaCl 6.8kcal 40.8kcal

Cu 3.11kcal 72.8kcal

Why is the heat of vaporization so much greater than the heat of fusion?

Energy Calculations
• PAY ATTENTION!!!!!!!!
• There are two separate types of calculations.
• 1) Heating/cooling calculations.
• 2) Phase Change calculations.
• DO NOT CONFUSE THESE TWO!!!!!!
Calculating
• How much heat does it take to completely vaporize 100. grams of room temperature water?
• The total heat required is the heat to raise the temperature of the water to its boiling point PLUS the heat to vaporize it at its boiling point.
Heating/Cooling of Materials
• Q = m * s * ΔT
• m is mass (in grams)
• s is specific heat capacity in J/(g ºC)
• ΔT is temperature change in K or ºC
Energy of a Phase Change
• H = n * ΔHpc
• n is in moles
• ΔH is in joules/mole or kcal/mol

How much heat does it take to completely vaporize 100. grams of room temperature(25 ºC) water?

• Q = 100. g * 4.18 J/g ºC * (100-25) ºC
• ΔHvap=(100. g/18.0 g/mol)*40.7 kJ/mol
• Total Heat Needed = Q + ΔH

= 31,400 J + 224,000 J = 255,000 J

Cooling
• How much heat is given off when 10.0 grams of water vapor at 100.0 ºC is converted into ice at –15.0 ºC ?
• s is 4.18 J/g ºC for liquid water
• s is 2.0 J/g ºC for water vapor
• s is 2.03 J/g ºC for solid water
Bonding in Metals
• Metals have loosely held electrons in their outer valence shell.
• These electrons can move from atom to atom to atom.
• This sharing creates the metallic bond – often described as atoms in a “sea of electrons”
The Metallic Bond
• The “Sea of Electrons” creates a strong but non-directional bond.
• Atoms in metals can move past one another when forced. This makes metals ductile.
• The motion of electrons makes metals electrically conductive.
Alloys
• Alloys are a material that has a mixture of elements and has metallic properties (conducts electricity and heat).
• Substitutional alloys are considered to be solid solutions.
Review
• What are the exothermic phase changes?
• Do these create a warming or cooling effect?
• What is happening to the potential energy?
Review
• What are the endothermic phase changes?
• Do these create a warming or cooling effect?
• What is happening to the potential energy?
Calculating
• How much energy would 1000. grams of ice at –10 o C absorb as it warms to room temperature?
• The total heat required is the heat to raise the temperature of the ice to its melting point PLUS the heat to melt it PLUS the heat to warm it to room temperature.