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Chapter 11 Slides

Chapter 11 Punchline

• Solids and Liquids are know as the “condensed states.” The properties of chemical species are best understood through the interactions between neighboring molecules (Intermolecular Forces).

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Part 1: The Phases of MatterCollege Chem I Definition

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Compressibility

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Phases of Matter

• Understand properties of the phases of matter on a molecular level.

• Phase is really about how the molecules are able to move– Degrees of Freedom

1. Translational– Gas and Liquid

2. Rotational– Gas and Liquid

3. Vibrational– Gas and Liquid and Solid

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Solids• The particles in a solid are packed close

together and are fixed in position

– though they may vibrate

• The close packing of the particles results in solids being incompressible

• The inability of the particles to move around results in solids retaining their shape and volume when placed in a new container, and prevents the solid from flowing

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Liquids• The particles in a liquid are closely

packed, but they have some ability to move around

• The close packing results in liquids being incompressible

• But the ability of the particles to move allows liquids to take the shape of their container and to flow – however, they don’t have enough freedom to escape or expand to fill the container

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Gases• Kinetic Molecular Theory• In the gas state, the particles have

complete freedom of motion and are not held together

• The particles are in constant motion, bumping into each other and the walls of the container

• There is a large amount of space between the particles– compared to the size of the particles– Low density– compressible

Competition of forces

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• Which phase a material is in depends largely on two major factors that are in competition with each other

1. the amount of kinetic energy the particles possess (Increases with increase in heat) (Forces the molecules apart)

2. the strength of attraction between the particles (Holds together)

(called Intermolecular forces)

1. Kinetic Energy

• Increasing the temperature increases the kinetic energy of the particles.

• The more kinetic energy the molecules have, the larger the intermolecular force they can overcome.

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2. Intermolecular Forces• The particles are attracted to each other by

electrostatic forces

• The strength of the attractive forces varies, depending on the molecule’s structure

• The stronger the attractive forces between the particles, the more they resist moving– Meaning, if the Intermolecular forces are strong, it

takes lots of kinetic energy to pull the molecules apart.

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Why Are Molecules Attracted to Each Other?

• Intermolecular attractions are due to attractive forces between opposite charges

• Larger charge = stronger attraction

• Longer distance = weaker attraction

• However, these attractive forces are small relative to the bonding forces between atoms– generally smaller charges

– generally over much larger distances

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Explain phases based on IMF and KE

• What do the movements of the molecules in these phases tell you about the competition between these forces? Which is “winning”?

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Explain phases based on IMF and KE

IMF

Gas

Solid

Liquid

Kinetic EnergyWhat do the movements of the molecules in these phases tell you about the competition between these forces? Which is “winning”?

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Phase Changes

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Causing a Phase Change with Pressure Change

• Kinetic–Molecular Theory of Gases assumes that when gas molecules collide, they have ZERO interaction with each other.

• This is usually a reasonable assumption because a “gas is like Kansas”– The distance between molecules is very large

• When you compress a gas, the space between the particles becomes much less, the intermolecular forces become larger.– Coulomb’s Law

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Liquefying a Gas

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Boiling Point and Melting Point

• The strength of the attractions between the particles of a substance determines its state

• The stronger the attractive forces are, the higher will be the boiling point of the liquid and the melting point of the solid

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IMF

IMF

Boiling Point

Part 4: Intermolecular Forces

• Ion-Dipole

• Hydrogen Bonding

• Dipole-Dipole

• Dispersion

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Same principle as ionic bonds, BUT much weaker than bonds:1. Distance between charges2. Smaller charges

Dispersion Forces• Fluctuations in the electron distribution in atoms and

molecules result in a temporary dipole

• The temporary dipole induces a dipole in surrounding molecules.

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The Noble gases are all

nonpolar atomic elements

As the molar mass increases,

the number of electrons

increases.

(The electron cloud gets

larger)

Effect of Molecular Sizeon Size of Dispersion Force

The stronger the attractive

forces between the molecules,

the higher the boiling point will

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Effect of Molecular Shapeon Size of Dispersion Force

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Dipole–Dipole Force• Polar molecules have a permanent dipole

• The permanent dipole adds to the attractive forces between the molecules– raising the boiling and melting points relative to nonpolar

molecules of similar size and shape

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Effect of Dipole–Dipole Attraction on Boiling and Melting Points

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H-Bonding

HF

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H-Bonding in Water

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Hydrogen Bonding (ice)

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The Bigger Picture:Guanine and Cytosine

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Ion-Dipole Forces

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Comparing Types of IMF

Note that the strength increase is relative. You must compare molecules of similar molecular mass and shape!

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Part 5: Properties of Matter that are affected by IMF

• Boiling Point/Melting Point

• Surface Tension

• Viscosity

• Capillary Action

• Vapor Pressure

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Boiling Point/Melting Point

• Stronger Intermolecular forces lead to higher melting and boiling points.

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or

Practice – Choose the substance in each pair with the higher boiling point

a)CH2FCH2F CH3CHF2

b)

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Practice – Choose the substance in each pair with the higher boiling point

more polara)CH2FCH2F CH3CHF2

b)

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or

polar nonpolar

a) CH3OH CH3CHF2

b) CH3-O-CH2CH3 CH3CH2CH2NH2

Practice – Choose the substance in each pair that is a liquid at room temperature (the other is a gas)

can H-bond

can H-bond

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Surface Tension

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Viscosity

Stronger IMF

More Viscous

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Capillary Action

Adhesive force vs. cohesive force

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Vapor Pressure

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Evaporation

When a molecule escapes into the gas phase, it takes its portion of the KE with it.

The rest of the KE is redistributed, but if the lost energy is not replaced by the surroundings, the overall temperature will decrease.

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Vapor Pressure with Lids

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Vapor Pressure and Temperature

Increasing the temperature increases the number of molecules with enough KE to escape the liquid phase, so the vapor pressure always increases.

Note that the trend is NOT linear.

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Clausius-Clapeyron Equation

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Part 6: Phase Diagrams and Phases of Matter

• A phase diagram is like a map that shows the stable phase at all pressure/temperature combinations

• The boundary lines indicate a combination of P and T where 2 phases exist in equilibrium

• Be able to draw and label a basic phase diagram and navigate one.

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Part 6: Phase Diagrams and Phases of Matter

Check out the fusion curve!

Although not common, water is actually more dense as a liquid than a solid, so the slope of the fusion curve has a negative slope.

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Supercritical fluid

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Heating Curves

• Allows you to calculate the heat energy put into or removed from a system• Allows you to calculate what temperature a sample is at after heating or

cooling.• When you encounter one of these problems, you should always DRAW the

heating curve… It makes things MUCH easier to see!

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Part 7: Crystalline Solids

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Part 7: Crystalline Solids

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Part 7: Crystalline Solids

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Examples of Crystal Structures

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Phase Diagram for Sulfur

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The above is more squished both in angles and lengths.

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