Lecture Myocardial Infarction

7/29/2019 Lecture Myocardial Infarction

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Announcements

Homework 27 and Homework 27 answers willbe available on Blackboard after lecture.

Quiz IX next time. Exam V will be returned at the end of lecture.

Your course-to-date percentage will beavailable on Blackboard under the heading

3rd% after lecture.


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Recap

The forces of attraction that give liquids and solids their fixed

volumes are called intermolecular forces.

Intermolecular forces are forces of attraction between

molecules or between molecules and ions.

There are four types of intermolecular forces, also known as

van der Waals forces:

Ion-dipole forces

Dipole-dipole forces

Hydrogen bonds

London dispersion forces

The relative strengths of the intermolecular forces:

Ion-dipole > H-bond > dipole-dipole > London dispersion

A substances boiling point is a measure of the strength of itsintermolecular forces.


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London Dispersion Forces London dispersion forces are the attractive forces between nonpolar

molecules.

CO2 or Br2

All atoms and molecules exhibit London dispersion forces. Because they arise from induced dipoles, London dispersion forces

are the weakest intermolecular forces.


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London Dispersion Force Strength

The relative strength of the London dispersion force

depends on thepolarizability of the molecule.

Polarizability is the ease with which a moleculeselectron cloud is distorted by a nearby electric field.

The greater the polarizability, the greater the London

dispersion force strength.

Polarizability correlates with molar mass.

London dispersion force polarizability molar mass


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Rank Br2, F2,

and I2according to

increasing

Londondispersion

force strength.

I2


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Ranking of Intermolecular Forces

The relative strengths of the intermolecular forces:

Ion-dipole > H-bond > dipole-dipole > London dispersion

A substances boiling point is a measure of the strength of its

intermolecular forces.

As intermolecular force strength increases, boiling point

increases.

This same trend is observed in melting points.

Formula Intermolecular Force Boiling Point (oC)

O2 London -183

HCl Dipole-dipole -85NH3 H-bond -33


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Rank the following compounds

according to increasing boiling point:CO2, H2O, and SO2.

CO2


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Liquids, Solids, and

Phase Changes


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Viscosity Viscosity is a measure of a liquids resistance to flow.

Van der Waals force strength: glycerol > water > ethanol >

benzene > pentane

Viscosity van der Waals force strength

Viscosity particle size


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Types of Solids

At their broadest level, solids can be classified

as either crystalline or amorphous.

Crystalline solids are those whose atoms, ions,or molecules have an ordered arrangementextending over a long range.

E.g., iron, table salt, ice, and diamond.

Amorphous solids are those whose constituent

parts are randomly arranged and have no long-range order.

Rubber


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Types of Crystalline Solids

Crystalline solids can be subclassified as:

Metallic

Ionic Molecular

Covalent network The different subtypes of crystalline solids

have different forces of attraction between the

particles.

Because they have different forces of

attraction, different subtypes of crystallinesolids have different physical properties.


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Metallic Solids

Metallic solids are the simplest examples of

crystalline solids because their constituent particles

are atoms, which can be approximated as spheres.

How can spheres be packed together in an ordered

arrangement that extends over a long range?

There are four typically observed results: Simple cubic packing

Body-centered cubic packing

Hexagonal closest packing

Cubic closest packing


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Unit Cells

Because these crystalline solids exhibit long-range

order, it isnt necessary to know what the extended

solid looks like.

Knowing the arrangement of the fundamental unit

that is repeated to compose the overall solid is

enough. The unit cell is the fundamental repeating unit that

makes up the overall solid.

There are three cubic unit cells:

Primitive cubic

Body-centered cubic Face-centered cubic


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Primitive and Body-Centered

Cubic Unit Cells


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Face-Centered Cubic Unit Cells


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Metallic Solids Summary There are two important variables that describe each packing type:

Number of nearest neighbors (aka coordination number): the number of other

atoms adjacent to any particular atom.

Packing efficiency: the percentage of unit cell volume occupied by atoms.

Conclusion: Particles pack together in solids as closely as possible,

maximizing their interparticle attractions.

Packing Type Unit Cell

Number of

Nearest

Neighbors

Packing

Efficiency

Number of

Metals with

Packing Type

Simple cubic Primitive cubic 6 52% 1

Body-centered

cubic

Body-centered

cubic8 68% 16

Hexagonalclosest

12 74% 21

Cubic closestFace-centered

cubic12 74% 18


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What Bonding in Metallic Solids Isnt

Metallic bonds cant be ionic.

Metals can readily lose electrons to satisfy the octet

rule, but they cant gain enough electrons to do so.Na+ is possible. Na7- is not.

Metallic bonds cant be localized or metals wouldnt

be malleable (capable of being beaten into shapes). At the atomic level, being beaten into new shapes

means moving atoms around, distorting the packing

arrangement.

If metallic bonds were localized, distorting the

packing arrangement would amount to breakingbonds, which requires an enormous amount of energy.


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What Bonding in Metallic Solids Is Metallic bonds involve delocalized sharing of electrons.

Molecular orbital theory is best at describing delocalized sharing of

electrons.

Increasing the number of interacting atoms increases the number of MOs. As the number of MOs increases, the energy differences between the MOs

decreaes.

(Conduction band)

(Valence band)


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Conductors, Insulators, and

Semiconductors In conductors such as metals, the conduction and valence bands overlap.

In insulators, the gap between the valence and conduction bands is too

great to be spanned.

In semiconductors, the gap between the valence and conduction bands is

small enough that it can be spanned by reasonably sized electric potentials.


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Ionic Solids

Like metallic solids, the particles that compose

ionic solids can be approximated as spheres.

Unlike metallic solids, the particles thatcompose ionic solids (cations and anions) have

charges and different sizes. Because of their charges, the nearest neighbors

of each cation will be anions, and vice versa.


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Sodium Chloride Unit Cell


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Bonding and Properties Ionic Solids

Ionic solids have high melting points because

melting somewhat overcomes the very strong

ionic bonds. Ionic solids are electrical nonconductors

because the ions are essentially fixed in place.

Ionic solids are hard because distorting the

solid requires moving ionsovercoming the

ionic bonds.

Ionic solids are brittle because when ions are

moved, strong repulsions result.


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Molecular Solids

Molecular solids are held together by van der

Waals forces.

Because van der Waals forces are typicallyweaker than metallic or ionic bonds, molecular

solids tend to be soft, easily deformed, and

possessed of low melting points.

Because there are no charged species and the

attractions are between molecules, molecular

solids are typically nonconductors.


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Ice Lattice Structure Solid water is held together by hydrogen bonds.

The hydrogen bonds cause water to have a hexagonal symmetry

because a hexagonal arrangement maximizes the hydrogen bonding

attractions.


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Covalent Network Solids

Covalent network solids are held together by

covalent bonds.

Since covalent bonds are strong and localized,covalent network solids tend to be hard, high

melting, electrical nonconductors.


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Graphite and Diamond Lattice

Structures


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Summary of Crystalline Solids and

Their Properties


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What inter-

particleforces

maintain

solid

krypton?

Ion-ionfo

rce

Ion-dipolef

orc

Hydrogenb

on

Dipole-dipolefor

Londondispers

ion.

Metallicb

on

Covalentbon

0% 0% 0% 0%0%0%0%

a. Ion-ion forces

b. Ion-dipole forces

c. Hydrogen bondsd. Dipole-dipole forces

e. London dispersion forces

f. Metallic bondsg. Covalent bonds


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Phase Changes

Phase changes, or changes of state, are changes in thephysical form of a substance that leave its chemical

identity intact.

E.g., boiling water converts liquid water into gaseous

water.

The names of the various phase changes: Fusion (melting): solid to liquid

Freezing: liquid to solid

Vaporization: liquid to gas Condensation: gas to liquid

Sublimation: solid to gas

Deposition: gas to solid


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Enthalpy, Entropy, & Phase Changes


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Energetics of Phase Changes

How much heat is required to convert 1.00 mol of H2O at -25.0 to 125.0 oC?

1. Heat ice from -25.0 oC to 0.0 oC.

q1 = nCH2O(s)T

q1 = (1.00 mol)(0.03657 kJ/molo

C)(0.0 25.0o

C) = 0.914 kJ2. Melt ice at 0.0 oC.

q2 = nHfusion = (1.00 mol)(6.01 kJ/mol) = 6.01 kJ

3. Heat liquid water from 0.0 oC to 100.0 oC.

q3 = nCH2O(l)T

q3 = (1.00 mol)(0.0754 kJ/moloC)(100.0 0.0 oC) = 7.54 kJ

4. Boil water at 100.0 oC.

q4 = nHvap = (1.00 mol)(40.67 kJ/mol) = 40.67 kJ

5. Heat steam from 100.0 oC to 125.0 oC.

q5 = nCH2O(g)T

q5 = (1.00 mol)(0.0331 kJ/moloC)(125.0 100.0 oC) = 0.828 kJqtotal = q1 + q2 + q3 + q4 + q5 = 55.96 kJ


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Heating Curve for 1 mol of H2O

l l l i f


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Molecular-Level Interpretation of

H2O Heating Curve Temperature changes when heating or cooling within a phase. Ergo,

particle kinetic energy (molecular speed) is changing.

Temperature is constant during a phase change. Thus, potential energy(intermolecular force) is changing.


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Evaporation Why will a glass of water left out on a table eventually evaporate

completely, without being heated?

At a given temperature, the molecules in the water have a distribution of

speeds, some of which are great enough to allow the molecules to escape

the liquid phase.


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Evaporation and Vapor Pressure


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Vapor Pressure

Vapor pressure (Pvap) is the pressure exerted bya vapor over its liquid.

What factors affectvapor pressure?

Vapor pressure T

Vapor pressure 1/van der Waals force strength


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Clausius-Clapeyron Equation The relationship between vapor pressure and temperature is given

by the Clausius-Clapeyron equation:

lnPvap = -(Hvap/R)(1/T) + C

The Clausisus-Clapeyron equation has the form of a straight line.

y = mx + b


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Clausius-Clapeyron Equation Example

When data is scarce, a two-point version of theClausius-Clapeyron equation may be used:

ln(P2/P1) = -(

Hvap/R)(1/T2 1/T1)


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ln(P2/P1) = -(Hvap/R)(1/T2 1/T1)

Butane lighters typically contain a mixture of liquid and

vapor. At 25 oC the vapor pressure of butane is 2.3 atm.

What is the pressure in the container at 150 oC? (The

Hvap of butane is 24.3 kJ/mol.)

2.90at

6.66at

18.1at

41.6at

0% 0%0%0%

a. 2.90 atm

b. 6.66 atmc. 18.1 atm

d. 41.6 atm


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Vapor Pressure and Boiling Point

The boiling point is the temperature at which the vaporpressure equals the external pressure.

The normal boiling pointis the temperature at which the

vapor pressure equals an external pressure of 1 atm.


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Summary

Viscosity van der Waals force strength; viscosity particle size

Most metallic solids exhibit hexagonal closest or cubic closest packing

because doing so packs the atoms as closely together as possible,

maximizing interparticle attractions.

Metallic bonds involve delocalized sharing of electrons, which is best

described as bands of MOs.

Ionic solids pack similarly to metals, except that the particles that

compose ionic solids (cations and anions) have charges and differentsizes.

Molecular solids are held together by van der Waals forces.

Covalent network solids are held together by covalent bonds. Phase changes, or changes of state, are changes in the physical form of

a substance that leave its chemical identity intact.

Vapor pressure is the pressure exerted by a vapor over its liquid.

The relationship between vapor pressure and temperature is given by

h l i l i l ( / )( / )

Lecture Myocardial Infarction

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