Intermolecular Forces © 2009, Prentice-Hall, Inc. A Molecular Comparison of Liquids and Solids The...

IntermolecularForces

© 2009, Prentice-Hall, Inc.

A Molecular Comparison of A Molecular Comparison of Liquids and SolidsLiquids and Solids

The fundamental difference between states of matter is the distance between particles.



States of MatterBecause in the solid and liquid states particles are closer together, we refer to them as condensed phases.



The States of Matter

• The state a substance is in at a particular temperature and pressure depends on two antagonistic entities:

– the kinetic energy of the particles;

– the strength of the attractions between the particles.



Intermolecular ForcesIntermolecular Forces

The attractions between molecules are not nearly as strong as the intramolecular attractions that hold compounds together.



Intermolecular Forces

They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities.



Intermolecular Forces

These intermolecular forces as a group are referred to as van der Waals forces.

• Dipole-dipole interactions• Hydrogen bonding• London dispersion forces



Ion-Dipole Interactions

• Ion-dipole interactions (a fourth type of force), are important in solutions of ions.

• The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents.



Dipole-Dipole Interactions

• Molecules that have permanent dipoles are attracted to each other.– The positive end of one is

attracted to the negative end of the other and vice-versa.

– These forces are only important when the molecules are close to each other.



Dipole-Dipole Interactions

The more polar the molecule, the higher is its boiling point.



London Dispersion Forces

While the electrons in the 1s orbital of helium would repel each other (and, therefore, tend to stay far away from each other), it does happen that they occasionally wind up on the same side of the atom.




At that instant, then, the helium atom is polar, with an excess of electrons on the left side and a shortage on the right side.




Another helium nearby, then, would have a dipole induced in it, as the electrons on the left side of helium atom 2 repel the electrons in the cloud on helium atom 1.




London dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and an induced dipole.




• These forces are present in all molecules, whether they are polar or nonpolar.

• The tendency of an electron cloud to distort in this way is called polarizability.



Factors Affecting London Forces

• The shape of the molecule affects the strength of dispersion forces: long, skinny molecules (like n-pentane tend to have stronger dispersion forces than short, fat ones (like neopentane).

• This is due to the increased surface area in n-pentane.



Factors Affecting London Forces

• The strength of dispersion forces tends to increase with increased molecular weight.

• Larger atoms have larger electron clouds which are easier to polarize.



Which Have a Greater Effect?Dipole-Dipole Interactions or Dispersion Forces

• If two molecules are of comparable size and shape, dipole-dipole interactions will likely the dominating force.

• If one molecule is much larger than another, dispersion forces will likely determine its physical properties.


Practice1. The dipole moments of acetonitrile, CH3CN, and methyl iodide,

CH3I, are 3.9 D and 1.62 D, respectively.

(a)Which of these substances has greater dipole–dipole attractions among its molecules?

(b) Which of these substances has greater London dispersion attractions?

(c) The boiling points of CH3CN and CH3I are 354.8 K and 315.6 K, respectively. Which substance has the greater overall attractive forces?

2. Of Br2, Ne, HCl, HBr, and N2, which is likely to have

(a) the largest intermolecular dispersion forces,

(b) the largest dipole–dipole attractive forces?




How Do We Explain This?

• The nonpolar series (SnH4 to CH4) follow the expected trend.

• The polar series follows the trend from H2Te through H2S, but water is quite an anomaly.



Hydrogen Bonding

• The dipole-dipole interactions experienced when H is bonded to N, O, or F are unusually strong.

• We call these interactions hydrogen bonds.



Hydrogen Bonding

• Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine.

Also, when hydrogen is bonded to one of those very electronegative elements, the hydrogen nucleus is exposed.


Practice1. In which of the following substances is hydrogen

bonding likely to play an important role in determining physical properties:

methane (CH4), hydrazine (H2NNH2), methyl fluoride (CH3F), or

hydrogen sulfide (H2S)?

2. In which of the following substances is significant hydrogen bonding possible:

methylene chloride (CH2Cl2), phosphine (PH3), hydrogen peroxide

(HOOH), or acetone (CH3COCH3)?




Summarizing Intermolecular Forces


More practice1. List the substances BaCl2, H2, CO, HF,

and Ne in order of increasing boiling points.

2. Given the following:

CH3CH3, CH3OH, and CH3CH2OH.

(a) Identify the intermolecular attractions present in the following substances

(b) select the substance with the highest boiling point:




Some Properties of LiquidsSome Properties of LiquidsThe strength of the attractions between particles can greatly affect the properties of a substance or solution.

With increasing IMF….

•Higher Viscosity

•Higher Surface Tension

•Higher Boiling and Melting Point

•Lower Equilibrium Vapor Pressure



Viscosity• Resistance of a liquid

to flow is called viscosity.

• It is related to the ease with which molecules can move past each other.

• Viscosity increases with stronger intermolecular forces and decreases with higher temperature.


Properties of Liquids

Cohesion is the intermolecular attraction between like molecules

11.3

Adhesion is an attraction between unlike molecules

Adhesion

Cohesion


Concave Meniscus of Water

Chemistry; The Science in Context; by Thomas R Gilbert, Rein V. Kirss, and Geoffrey Davies, Norton Publisher, 2004, p 458



Surface Tension

Surface tension results from the net inward force experienced by the molecules on the surface of a liquid.



Phase ChangesPhase Changes



Energy Changes Associated with Changes of State

The heat of fusion is the energy required to change a solid at its melting point to a liquid.




The heat of vaporization is defined as the energy required to change a liquid at its boiling point to a gas.




• The heat added to the system at the melting and boiling points goes into pulling the molecules farther apart from each other.

• The temperature of the substance does not rise during a phase change.


Practice1. Calculate the enthalpy change upon converting

1.00 mol of ice at –25 °C to water vapor (steam) at 125 °C under a constant pressure of 1 atm. The specific heats of ice, water, and steam are 2.03 J/g-K, 4.18 J/g-K, and 1.84 J/g-K, respectively. For H2O, ΔHfus = 6.01 kJ/mol and ΔHvap = 40.67 kJ/mol.

2. What is the enthalpy change during the process in which 100.0 g of water at 50.0 °C is cooled to ice at –30.0 °C ?




Vapor Pressure

• At any temperature some molecules in a liquid have enough energy to escape.

• As the temperature rises, the fraction of molecules that have enough energy to escape increases.



Vapor Pressure

As more molecules escape the liquid, the pressure they exert increases.



Vapor Pressure

The liquid and vapor reach a state of dynamic equilibrium: liquid molecules evaporate and vapor molecules condense at the same rate.



Vapor Pressure• The boiling point of a

liquid is the temperature at which it’s vapor pressure equals atmospheric pressure.

• The normal boiling point is the temperature at which its vapor pressure is 760 torr.

* Estimate the boiling point of diethyl ether under an external pressure of 0.80 atm.* At what external pressure will ethanol have a boiling point of 60 °C?



Phase DiagramsPhase Diagrams(prior knowledge)(prior knowledge)

Phase diagrams display the state of a substance at various pressures and temperatures and the places where equilibria exist between phases.



Phase Diagrams

• The circled line is the liquid-vapor interface.• It starts at the triple point (T), the point at

which all three states are in equilibrium.



Phase Diagrams

It ends at the critical point (C); above this critical temperature and critical pressure the liquid and vapor are indistinguishable from each other.



Phase Diagrams

Each point along this line is the boiling point of the substance at that pressure.



Phase Diagrams

• The circled line in the diagram below is the interface between liquid and solid.

• The melting point at each pressure can be found along this line.



Phase Diagrams• Below the triple point the substance cannot

exist in the liquid state.• Along the circled line the solid and gas

phases are in equilibrium; the sublimation point at each pressure is along this line.



Phase Diagram of Water

• Note the high critical temperature and critical pressure.– These are due to the

strong van der Waals forces between water molecules.



Phase Diagram of Water

• The slope of the solid-liquid line is negative.– This means that as the

pressure is increased at a temperature just below the melting point, water goes from a solid to a liquid.



Phase Diagram of Carbon Dioxide

Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm; CO2 sublimes at normal pressures.



Structures of SolidsStructures of Solids

• We can think of solids as falling into two groups:– crystalline, in which

particles are in highly ordered arrangement.



Solids

• We can think of solids as falling into two groups:– amorphous, in which

there is no particular order in the arrangement of particles.



Attractions in Ionic CrystalsIn ionic crystals, ions pack themselves so as to maximize the attractions and minimize repulsions between the ions.



Bonding in SolidsBonding in Solids



Covalent-Network andMolecular Solids

• Diamonds are an example of a covalent-network solid, in which atoms are covalently bonded to each other.– They tend to be hard and have high melting

points.



Covalent-Network andMolecular Solids

• Graphite is an example of a molecular solid, in which atoms are held together with van der Waals forces.– They tend to be softer and have lower melting

points.


Metallic Bonding, Alloys & Semiconductors

Presented By, Mark Langella, APSI

Chemistry 2014 , PWISTA .com



Metallic Solids

• Metals are not covalently bonded, but the attractions between atoms are too strong to be van der Waals forces.

• In metals valence electrons are delocalized throughout the solid.


Alloys

• Alloys contain more than one element and have the characteristic properties of metals.

• Solid Solution alloys are homogeneous mixtures.

• Heterogeneous alloys: The components are not dispersed uniformly (e.g., pearlite steel has two phases: almost pure Fe and cementite, Fe3C).

• Pure metals and alloys have different physical properties.

• An alloy of gold and copper is used in jewelry (the alloy is harder than the relatively soft pure 24 karat gold).

• 14 karat gold is an alloy containing 58% gold.




Metal Alloys-Solid SolutionsSubstance has mixture of

element and metallic properties.

1.Substitutional Alloy: some metal atoms replaced by others of similar size. Electronegativities usually are similar. The atoms must have similar atomic radii. The elements must have similar bonding characteristics.

•brass = Cu/Zn




Metal Alloys(continued)

2.Interstitial Alloy: Interstices (holes) in closest packed metal structure are occupied by small atoms. Solute atoms occupy interstices “small holes” between solvent atoms. One element (usually a nonmetal) must have a significantly smaller radius than the other (in order to fit into the interstitial site).

steel = iron + carbon

3.Both types: Alloy steels contain a mix of substitutional (Cr, Mo) and interstitial (Carbon) alloys.




Substitutional Alloy

Interstitial Alloy




Alloys vs. Pure Metal

• The alloy is much harder, stronger, and less ductile than the pure metal (increased bonding between nonmetal and metal).

An example is steel (contains up to 3% carbon).

mild steels (<0.2% carbon) - useful for chains, nails, etc.

medium steels (0.2-0.6% carbon) - useful for girders, rails, etc.

high-carbon steels (0.6-1.5% carbon) - used in cutlery, tools, springs.

Other elements may also be added to make alloy steels.

Addition of V and Cr increases the strength of the steel and improves its resistance to stress and corrosion.

The most important iron alloy is stainless steel. It contains C, Cr (from ferrochrome, FeCr2), and Ni.







Which two substances are most likely to form an interstitial alloy?

• Nickel and titanium

• Silver and tin

• Tin and lead

• Copper and zinc

• Tungsten and carbon




Which two substances are most likely to form an interstitial alloy?

• Nickel and titanium

• Silver and tin

• Tin and lead

• Copper and zinc

• Tungsten and carbon




Bonding Models for Metals

APSI 2014 PWISTA.com


Bonding Models for Metals•Electron Sea Model: A regular array of metals in a “sea” of electrons. The electron-sea model is a qualitative interpretation of band theory (molecular-orbital model for metals).•Band (Molecular Orbital) Model: Electrons assumed to travel around metal crystal in MOs formed from valence atomic orbitals of metal atoms.•Conduction Bands: closely spaced empty molecular orbitals allow conductivity of heat and electricity.




Molecular Orbital Theory

Recall that atomic orbitals mix to give rise to molecular orbitals.


Molecular-Orbital Model for Metals

• Delocalized bonding requires the atomic orbitals on one atom to interact with atomic orbitals on neighboring atoms.

• Example: Graphite electrons are delocalized over a whole plane, while benzene molecules have electrons delocalized over a ring.

• Recall that the number of molecular orbitals is equal to the number of atomic orbitals. Each orbital can hold two electrons.

• In metals there are a very large number of orbitals.• As the number of orbitals increases, their energy spacing

decreases and they band together.• The available electrons do not completely fill the band of

orbitals.





Molecular-Orbital Model for Metals

• Therefore, electrons can be promoted to unoccupied energy bands.

• Because the energy differences between orbitals are small the promotion of electrons requires little energy.

• As we move across the transition metal series, the antibonding band starts becoming filled.

• Therefore, the first half of the transition metal series has only bonding-bonding interactions and the second half has bonding–antibonding interactions.

• We expect the metals in the middle of the transition metal series (group 6B) to have the highest melting points.

• The energy gap between bands is called the band gap.


IntermolecularForcesAPSI 2014 PWISTA.com

Molecular Orbital Theory

In such elements, the energy gap between molecular orbitals essentially disappears, and continuous bands of energy states result.


Formation of Bands


When atoms come together to form a compound, their atom orbital energies mix to form molecular orbital energies. As more atoms begin to mix and more molecular orbitals are formed, it is expected that many of these energy levels will start to be very close to, or even completely degenerate, in energy. These energy levels are then said to form bands of energy remember each orbital only holds two electrons.


The electronic band structure of nickel.

The left side of the figure shows the electron configuration of a single Ni atom, while the right-hand side of the figure shows how these orbital energy levels broaden into energy bands in bulk nickel. The horizontal dashed gray line denotes the position of the Fermi Level, which separates the occupied molecular orbitals (shaded in blue) from the unoccupied molecular orbitals.



Presented By, Mark Langella, APSI Chemistry

2014 , PWISTA .com

Types of MaterialsRather than having molecular orbitals separated by an energy gap, these substances have energy bands.

The gap between bands determines whether a substance is a metal, a semiconductor, or an insulator.


Energy bands in metals, semiconductors, and insulators.

Metals are characterized by the highest-energy electrons occupying a partially filled band. Semiconductors and insulators have an energy gap that separates the completely filled band (shaded in blue) and the empty band (unshaded), known as the band gap and represented by the symbol Eg. The filled band is called the valence band (VB), and the empty band is called the conduction band (CB). Semiconductors have a smaller band gap than insulators.



IntermolecularForcesPresented By, Mark

Langella, APSI Chemistry 2014 , PWISTA .com

Metals

• Valence electrons are in a partially-filled band.

• There is virtually no energy needed for an electron to go from the lower, occupied part of the band to the higher, unoccupied part.

• This is how a metal conducts electricity.

IntermolecularForcesPresented By, Mark

Langella, APSI Chemistry 2014 , PWISTA .com

Insulators

• The energy band gap in insulating materials is generally greater than ~350 kJ/mol.

• They are not conductive.



2014 , PWISTA .com

SemiconductorsSemiconductors have a gap between the valence band and conduction band of ~50-300 kJ/mol.


An intrinsic semiconductor is a semiconductor in its pure state. For every electron that jumps into the conduction band, the

missing electron will generate a hole that can move freely in the valence band. The number of holes will equal the number of

electrons that have jumped. The higher the temp more electrons into conduction band.




The following pictures show the electron populations of the bands of MO energy levels for four different materials:

Classify each material as an insulator, a semiconductor, or a metal.

Arrange the four materials in order of increasing electrical conductivity. Explain.

Tell whether the conductivity of each material increases or decreases when the temperature increases.

(a)

(b)

(c)Presented By, Mark

Langella, APSI Chemistry 2014 ,

PWISTA .com



2014 , PWISTA .com

Semiconductors• Among elements, only silicon,

germanium and graphite (carbon), all of which have 4 valence electrons, are semiconductors.

• Inorganic semiconductors (like GaAs) tend to have an average of 4 valence electrons (3 for Ga, 5 for As).



2014 , PWISTA .com

Doping

By introducing very small amounts of impurities that have more valence electrons (n-Type) or fewer (p-Type) valence electrons, one can increase or decrease the conductivity of a semiconductor.


The addition of controlled small amounts of impurities (doping) to a semiconductor changes the electronic properties of the material.

•Left: A pure, intrinsic semiconductor has a filled valence band and an empty conduction band (ideally). •Middle: The addition of a dopant atom that has more valence electrons than the host atom adds electrons to the conduction band (i.e., phosphorus doped into silicon). The resulting material is an n-type semiconductor. •Right: The addition of a dopant atom that has fewer valence electrons than the host atom leads to fewer electrons in the valence band or more holes in the valence band (i.e., aluminum doped into silicon). The resulting material is a p-type semiconductor.
















Which of the following is a p-type semiconductor?

• Sulfur-doped carbon• Boron-doped germanium • Phosphorus-doped silicon• Ultra-pure silicon• Carbon-doped copper




Which of the following is a p-type semiconductor?

• Sulfur-doped carbon• Boron-doped germanium • Phosphorus-doped silicon• Ultra-pure silicon• Carbon-doped copper




Which of the following elements, if doped into silicon, would yield an n-type semiconductor? Ga; As; C.

Suggest an element that could be used to dope silicon to yield a p-type material.

Practice Exercises




Diode- Used to switch and convert between electromagnetic radiation and

electric current

• Semiconductor created that has p-type on one half and n-type on the other half

• Known as “p-n junction”







Light emitting diodes.

The heart of a light emitting diode is a p-n junction where an applied voltage drives electrons and holes to meet. Bottom: The color of light emitted depends upon the band gap of the semiconductor used to form the p-n junction. For display technology red, green, and blue are the most important colors because all other colors can be made by mixing these colors.




Presented By, Mark Langella, APSI Chemistry 2014 , PWISTA .com

Color λ Voltage Drop Composition

Red 610 < λ < 760 1.63 < ΔV < 2.03

Aluminium gallium arsenide (AlGaAs)Gallium arsenide phosphide (GaAsP)Aluminium gallium indium phosphide (AlGaInP)Gallium(III) phosphide (GaP)

Orange 590 < λ < 610 2.03 < ΔV < 2.10Gallium arsenide phosphide (GaAsP)Aluminium gallium indium phosphide (AlGaInP)Gallium(III) phosphide (GaP)

Yellow 570 < λ < 590 2.10 < ΔV < 2.18Gallium arsenide phosphide (GaAsP)Aluminium gallium indium phosphide (AlGaInP)Gallium(III) phosphide (GaP)

Green 500 < λ < 570 1.9[63] < ΔV < 4.0

Traditional green:Gallium(III) phosphide (GaP)Aluminium gallium indium phosphide (AlGaInP)Aluminium gallium phosphide (AlGaP)Pure green:Indium gallium nitride (InGaN) / Gallium(III) nitride (GaN)

Blue 450 < λ < 500 2.48 < ΔV < 3.7

Zinc selenide (ZnSe)Indium gallium nitride (InGaN)Silicon carbide (SiC) as substrateSilicon (Si) as substrate—under development

Violet 400 < λ < 450 2.76 < ΔV < 4.0 Indium gallium nitride (InGaN)


Solar Cells



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