chapter 03_liquid,solid n phase changes

8/14/2019 Chapter 03_liquid,Solid n Phase Changes

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Chapter 3:Liquids, Solids & PhaseChanges

Chapter Objectives:

To learn the differences between the solid, liquid, and gas state,

and how the polarity of molecules influences those states.

To learn the different types of intermolecular forces between

different molecules.


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Kinetic Molecular Theory of Liquidsand Solids

Liquids and solids have significant interactions.

Liquids and solids have well-defined volume.

Liquid molecules flow, while solids are held rigid.

A molecular comparison of gases, liquids, and solids.

(b) In gases, the particles feel little attraction for one another

and are free to move about randomly.

(b) In liquids, the particles are held close together by attractive

forces but are free to move over one another.

(c) In solids, the particles are rigidly held in an ordered arrangement.


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Polar Bonds Individual covalent bonds are polar if the atoms beingconnected are of different electronegativities. This is

described as a bond dipole.

Example: CH3Cl

The CH bonds are nonpolarsince C and H have about thesame electronegativity.

Since Cl is more electronegative than C, the CCl bond is

polarizedso that the Cl atom is slightly electron-rich (partial

negative charge,-) and the C atom is slightly electron-poor

(partial positive charge, +). This bond is a polar covalent

bond (or just polar bond).


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Dipole Moment (): the sum of all the bond dipoles within a molecule,

(net polarity) can be illustrated with an electrostatic potential map.

Since CH3Cl has a tetrahedralshape, with one polar bond and three nonpolar

bonds, there is an overall dipole moment pointing towards the Cl end of themolecule.

Depending on the number and orientation of the bond dipoles, the molecule may

possess an overall molecular dipole.

Polar Bonds


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Ammonia and water

Carbon dioxide and tetrachloromethaneStructures and dipole moments for ammonia and water are shown

Structures and dipole moments are shown for carbon dioxide and

tetrachloromethane. Although each molecule has bond dipoles,

they do not have molecular dipoles


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Intramolecular &Intermolecular Forces

Intramolecular forcesoperate withineach molecule,influencing the chemicalproperties of the substance

(i.e., covalent bonds).

These are the forces that hold the atoms in a molecule together.

They are very strong forces which result from large charges (onprotons and electrons)interacting over very short distances.

Intermolecular forces (van der Waals forces) operate

betweenseparate molecules, influencing thephysical

properties of the substance.

These are the forces that hold liquids and solids together, and

influence their melting and boiling points. They are weaker forces,

because they result from smaller charges, or partial charges,

interacting over much larger distances.


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To break an OH bond in water, the water must be

heated to thousands of degrees C; to completely

overcome the intermolecular forces, all you have to do is

boil it 100C.



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(b) At a higher temperature,

intermolecular forces are no longer

able to keep molecules close

together, so nitrogen becomes a

gas.


(a) In an individual N2 molecule,atoms are held together by strong

intramolecular force(covalent bond).

Different N2 molecules are weakly

attracted to one another at low

temperature by intermolecular forces,

causing nitrogen to become liquid.


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An attractive interaction between molecules

Determine bulk properties of matter.

Much weaker than intramolecular forces

Several types of intermolecular (IM) forces:

Iondipole

DipoledipoleLondon dispersion forces

Hydrogen bonds.

Intermolecular Forces


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London Dispersion

Forces Attraction is due to instantaneous, temporary dipoles

formed due to electron motions

The larger the molar mass of a molecule, generally the

greater the LDFs.

Dispersion forces exist between allmolecules, but they

are the onlyforces that exist between nonpolar

molecules.


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Using molecular bromine as an example, each molecule has zero

polarity.However, due to the motion of electrons at any given instant a

temporary dipole would arise that would then induce a dipole in an

adjacent molecule.This type of intermolecular attraction is called London dispersion

forces.

London Dispersion

Forces


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Dipole-Dipole Forces Dipole-Dipole forces are the attractions between the oppositepartial charges in the permanent dipoles of polar molecules.

exit between all polar molecules, (HCl HCl)

(a) Polar molecules attract one another and approach

closely

when oriented with unlike charges together, but (b) they

repel one another and push apart when oriented with like

charges together.


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In general, for molecules of the same molecular

weight, a polar molecule (dipole-dipole +

London) will have a higher boiling point than a

nonpolar molecule (London only):

Dipole-Dipole Forces


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Attractive interaction between a hydrogen atombonded to a very electronegative atom (O, N, or

F) and an unshared electron pair on another

electronegative atom.

Hydrogen bonds are also found between molecules of water and

molecules of ammonia.

Hydrogen Bonding


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Ion-Dipole forces are the result of electrical interactionsbetween an ion and the partial charges on a polar

molecule.

Polar molecules orient toward ions so that

(a) the positive end of the dipole is near an anion and

(b) the negative end of the dipole is near a cation.

Ion-Dipole Forces


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These forces are responsible for the ability of

polar solvents like water to dissolve ionic

compounds.

Ion-Dipole Forces


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Ion- and Dipole-Induced DipoleForces

Ion - induced dipole forcesare the attractive forces that existbetween ions and nonpolar molecules.

Being next to an ion induces a dipole in a nonpolar molecule,attracting it towards the ion.

These forces are responsible for the attraction betweenFe2+ and O2 molecules in the bloodstream, and contributesto the solvation of ions in water.

Dipole - induced dipole forcesare the attractive forces that exist

between polar molecules and nonpolar molecules. Being next to apolar molecule induces a dipole in a nonpolar molecule.

These forces are responsible for the solvation of gases(nonpolar) in water (polar).


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Summary: Intermolecular

Forces

CH4 and CH4; F2

and F2; CH4 and

F2

two nonpolarmolecules

London(dispersion)

forces

HCl and Cl2a polar moleculeand a nonpolar

molecule

Dipole -Induced

dipole

Fe2+ and O2an ion and anonpolar molecule

Ion - Induceddipole

CH3Br and ICl;

CH3Br and H2O

two polar moleculesDipole - Dipole

H2O and H2O;

H2O

and CH3CH2OH

molecules whichhave H

on N, O, or F atoms

Hydrogen bond

Na+ and H2Oan ion and a polarmolecule

Ion-dipole

ExamplesFormed by theattractionbetween

Intermolecular Forces

Str

ength


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Intermolecular Force

Effects Surface Tension The resistance of a liquid to spread out andincrease its surface area.

Surface tension results from intermolecular

force differences between molecules in theinterior of a liquid and those on the surface.

Molecules at the surface of a liquid feel

attractive forces only one side and arethus pulled in toward the liquid, while

molecules in interior are surrounded and

are pulled equally in all direction.


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Surface tension, which causes these drops of liquid mercury to formbeads


Effects

Atoms on the surface are less stable because they have fewer neighbors

and feel fewer attractive forces than atoms in the interior, so the liquid acts

to minimize their number by minimizing the area of the surface.


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More surface tension examples


Effects


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Effects2. Viscosity The measure of a liquids resistance to flow and is related to the

ease with which molecules move around, and thus to the

intermolecular forces.

Substances composed of small, nonpolar molecules (such as

gasoline and benzene) have low viscosities.

Polar molecules (such as glycerol) and molecules composed of

long chains of atoms (such as oil and grease) have higherviscosities.

The viscosity of a liquid decreases at higher temperatures.


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Effects


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Phase Changes Enthalpy the heat flow associated with making or breaking

intermolecular attractions that hold liquids and solids together Entropy change in a molecular randomness between various phases

Changes from a less random phase to a more random one have positive

values of H and S. Changes from a more random phase to a less

random one have negative values of H and S.


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Sublimation:The process in which molecules go

directly from the solid into the vapor phase.

Deposition:The process in which molecules go

directly from the vapor into the solid phase.

Molar heat of sublimation (Hsub):The energy (kJ)

required to sublime one mole of solid.

Hsub=Hfus + Hvap

Phase Changes


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Phase Changes

Molar Heat of Fusion (Hfus):

The energy required to melt one mole of solid (in kJ).

Molar Heat of Vaporization (Hvap):

The energy (in kJ) required to vaporize one mole of

liquid.


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Heating curve for water

A heating curve for H2O, showing the temperature changes and phasetransitions that occur when heat is added. The plateau at 0C represents

the melting of solid ice, and the plateau at 100C represents the boiling

of liquid water.

Plateau regions in a heating curve indicate a change in phase of the

substance where the temperature remains constant


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Vapor Pressure: The pressure exerted by gaseous

molecules above a liquid.

Liquids after sitting for a length of time in (a) an open container

and (b) a closed container. The liquid in the open container hasevaporated, but the liquid in the closed container has brought

about a rise in pressure.

Evaporation of a liquid results in more gas phase molecules

which exert a pressure in a closed container

Phase Changes


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Phase changes depend on temperature.

The distribution of molecular kinetic energies in a liquid at two temperatures.

Only the faster-moving molecules have sufficient kinetic energy to escape

from the liquid and enter the vapor. The higher the temperature, the larger

the number of molecules with enough energy to escape.


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Kinds of Solids Solids are divided into two categories:

Crystalline:

- Possesses rigid and long-range order

- Flat faces

- Distinct angles

- eg. NaCl

Amorphous

- Lacks well-defined arrangement (particles are randomly

arranged)

- have no long ranged structure- eg. rubber

Structure of a crystalline solid is based on the unit cell, a basic

repeating structural unit.


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Crystalline Solids


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Crystal structures of (a) ice, a molecular solid, and (b) quartz, a covalentnetwork solid.

Ice consists of individual H2O molecules held together in a regularmanner by hydrogen bonds.

Quartz (SiO2) is essentially one very large molecule whose Si and Oatoms are linked by covalent bonds.

Each silicon atom has tetrahedral geometry and is bonded to fouroxygens; each oxygen has approximately linear geometry and is bondedto two silicons.

The shorthand representation on the right shows how SiO4 tetrahedrajoin at their corners to share oxygen atoms.


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Cubic Packing

Simple Cube and Body-Centered Cube:Simple Cube and Body-Centered Cube:

(a) Simple cubic packing of spheres

all the layers are identical and all atoms are lined up in stacks and rows.

(b) Body-centered cubic packing of spheres

spheres in layer a are separated slightly and the spheres in layer b

are offset so that they fit into the depressions between atoms in layer a.


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Hexagonal and Cubic Closest-

Packinga) Hexagonal closest-packing.

Two alternating hexagonal layers

(a and b) offset from each other so

that the spheres in one layer sit in

the small triangular depressions of

neighboring layers.

i) Cubic closest-packing of spheres

Three alternating hexagonal

layers, (a, b, and c) offset fromone

another so that the spheres in one

layer sit in small triangular

depressions of neighboring layers.


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Cubic Unit CellsSimple Cube and Body-Centered Cube:Simple Cube and Body-Centered Cube:

Geometries of (a) primitive-cubic and (b) body-centered cubic unit cells in

both a skeletal view (top) and a space-filling view (bottom).

Part (c) shows how eight primitive-cubic unit cells stack together to share

a common corner where they meet.


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Cubic Unit Cells

(a) Geometry of a face-

centered cubic unit cell,

(b) a view showing how thisunit cell is found incubic closest-packing.The faces are tilted at54.7 angles to thethree repeating atomiclayers

Face-Centered Cube:Face-Centered Cube:


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A Phase Diagram is a graphical display of thetemperatures and pressures at which twophases of a substance are in equilibrium.

Triple Point:The only condition under which all threephases can be in equilibrium with one another.

Critical Temperature (Tc):The temperature above which

the gas phase cannot be made to liquefy at any pressure.

Critical Pressure (Pc):The minimum pressure required to

liquefy a gas at its critical temp.

Phase Diagrams


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Phase diagram for water

A phase diagram for H2O, showing a negative slope for the solid/liquid

boundary.


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Phase diagram for

carbon dioxide

A phase diagram for CO2, showing a positive slope for

the solid/liquid boundary. The pressure and temperature

axes are not to scale.


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Once the critical temperature and pressure have been reached the

two distinct phases of liquid and gas are no longer visible. The

meniscus can no longer be seen. One homogenous phase called

the "supercritical fluid" phase occurs which shows properties of

b h li id d

Here we can see the separate phases of carbon dioxide. The meniscus

is easily observed.

With an increase in temperature the meniscus begins to diminish.

Increasing the temperature further causes the gas and liquid

densities to become more similar. The meniscus is less easily

observed but still evident.

Formation of supercritical fluid

chapter 03_liquid,solid n phase changes

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