chapters 10: intermolecular forces, liquids, and solids

IntermolecularForces

Chapters 10:Intermolecular Forces,

Liquids, and Solids

AP Chemistry – Unit 6


States of MatterThe 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 Forces

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



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



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


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 be the dominating force.

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


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.


Summarizing Intermolecular Forces


Example 10.1Arrange the following substances in the expected order of increasing boiling point: carbon tetrabromide, CBr4; butane, CH3CH2CH2CH3; fluorine, F2; acetaldehyde, CH3CHO.

Example 10.2In which of these substances is hydrogen bonding an important intermolecular force: N2, HI, HF, CH3CHO, and CH3OH? Explain.


Intermolecular Forces Affect Many Physical Properties

The strength of the attractions between particles can greatly affect the properties of a substance or solution.


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.


Surface Tension

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


Adhesive and Cohesive Forces

The liquid spreads, because adhesive forces

are comparable in strength to cohesive

forces.

The liquid “beads up.” Which forces are

stronger, adhesive or cohesive?


Meniscus Formation

Water wets the glass (adhesive forces) and its

attraction for glass forms a concave-

up surface.

What conclusion can we draw about the cohesive forces

in mercury?


Capillary Action

The adhesive forces wet more and more of the inside of the tiny tube, drawing water farther up

the tube.


Phase 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.



• Q = mcT is used for all segments of the graph where temperature is increasing.

• Q = mHf is used for the BC segment.

• Q = mHv is used for the DE segment.


Example 10.3Calculate the thermal energy required to convert 14 kg of ice at –12C to steam at 105C (heat capacity for ice, water and steam are 2.1 J/g.C, 4.2 J/g.C, 2.0 J/g.C respectively and Hv = 40.7 kJ/mol and Hf = 6.02 kJ/mol)


Cooling Curve for Water

The liquid water cools until …

… the freezing point is reached, at which time the

temperature remains constant

as solid forms.

Once all of the liquid has

solidified, the temperature again drops.

If the liquid is cooled carefully, it can supercool.


Vapor Pressure

• The vapor pressure of a liquid is the partial pressure exerted by the vapor when it is in dynamic equilibrium with the liquid at a constant temperature.

vaporization

Liquid Vaporcondensation


Liquid–Vapor Equilibrium

More vapor forms; rate of condensation of that

vapor increases …

… until equilibrium is attained.


Vapor pressure increases with temperature; why?


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 Curves

Which of the five compounds has the

strongest intermolecular forces?

How can you tell?

CS2

CH3OH

CH3CH2OH

H2O

C6H5NH2

What is the vapor pressure of H2O at 100 °C, according to this graph? What is the significance of that numeric value of vapor pressure?


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.


Phase Diagrams

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.


Phase Diagram for HgI2

HgI2 has two solid phases, red

and yellow.

As the vessel is allowed to cool, will the contents appear more red

or more yellow?


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.


Types of Crystalline Solids

– Molecular solids, containing molecules held to one another by dispersion/dipole–dipole/ hydrogen bonding forces

– Network (covalent) solids– Ionic solids– Metallic solids (metals)


Crystal Lattices

• To describe crystals, three-dimensional views must be used.

• The repeating unit of the lattice is called the unit cell.

• There are a number of different types of unit cell; hexagonal, rhombic, cubic, etc.

• The three types of cubic unit cells are: simple cubic, body-centered cubic (bcc), face-centered cubic (fcc).


Cubic Unit CellsThe unit cell is a cube in each

case.

Whole atoms shown for

clarity.


Close-Packed StructuresA close packed structure in two

dimensions.


Close Packing in 3 Dimensions

Two layers, stacked, give two different locations for the

third layer …

Third layer directly above first layer: HCP

Third layer over the octahedral holes in the second layer: CCP


Example 11.10Copper crystallizes in the cubic close-packed arrangement. The atomic (metallic) radius of a Cu atom is 127.8 pm. (a) What is the length, in picometers, of the unit cell in a sample of copper metal? (b) What is the volume of that unit cell, in cubic centimeters? (c) How many atoms belong to the unit cell?

Example 11.11

Use the results of Example 11.10, the molar mass of copper, and Avogadro’s number to calculate the density of metallic copper.


Ionic Crystal Structures

• Ionic crystals have two different types of structural units—cations and anions.

• The cations and anions ordinarily are different sizes.• Smaller cations can fill the voids between the larger

anions.• Where the cations go depends on the size of the

cations and on the size of the voids.• The smallest voids are the tetrahedral holes, then the

octahedral holes, and finally the holes in a cubic structure. Therefore …


Ionic Crystal Structures

• Tetrahedral hole filling occurs when the cations are small; when the radii ratio is:

0.225 < rc/ra < 0.414

• Octahedral filling occurs with larger cations, when the radii ratio is:

0.414 < rc/ra < 0.732

• The arrangement is cubic if rc/ra > 0.732.


Crystalline Solids

We can determine the empirical formula of an ionic solid by determining how many ions of each element fall within the unit cell.


Unit Cell of Cesium Chloride

How many cesium ions are inside this

unit cell? How many chloride ions?


Unit Cell of Sodium Chloride

How many sodium ions are inside this

unit cell? How many chloride ions?


Experimental Determinationof Crystal Structures

X rays

Angle of diffraction can be used to find distance d,

using simple trigonometry.


Network Covalent Solids• Network solids have a network of covalent bonds that

extend throughout the solid, holding it firmly together.• The allotropes of carbon provide good examples.

– Diamond has each carbon bonded to four other carbons in a tetrahedral arrangement using sp3 hybridization.

– Graphite has each carbon bonded to three other carbons in the same plane using sp2 hybridization.

– Fullerenes are roughly spherical collections of carbon atoms in the shape of a soccer ball.

– A nanotube can be thought of as a plane of graphite rolled into a tube.


Crystal Structure of Diamond

Three-dimensional network is extremely

strong, rigid.What kind of forces must be overcome to

melt diamond?


Crystal Structure of Graphite

Forces between layers are

relatively weak.

Hexagons of sp2-hybridized

carbon atoms.


Structure of a Buckyball

C60 molecule


Structure of a Nanotube

A nanotube can be thought of as a sheet of graphite,

rolled into a tube, capped with half of a buckyball.


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.

chapters 10: intermolecular forces, liquids, and solids

Documents

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stronger dispersion

induced dipole

instantaneous dipole

dispersion forcesif

london dispersion forceswhile

london dispersion forcesat

helium atom