Chapter 11 Slides - UCAfaculty.uca.edu/kdooley/su2014_chem1451_chapter_11_slides.pdfChapter 11...
Transcript of Chapter 11 Slides - UCAfaculty.uca.edu/kdooley/su2014_chem1451_chapter_11_slides.pdfChapter 11...
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
6Tro: Chemistry: A Molecular Approach, 2/e
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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
Tro: Chemistry: A Molecular Approach, 2/e
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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
be.21
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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