[Session 9]

[Intermolecular Attractions&

Properties of Liquids and Solids]

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Learning Outcomes

Students are able to understand about

intermolecular attractions and determine

properties in liquid and solid that are

affected by intermolecular attractions

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Outline

• Types of Intermolecular Forces

• Properties that depend on the tightness of packing

• Properties that depend on intermolecular attractions

• Equilibria in state changes

• Energy and changes of state

• Le Chatelier’s Principle

• Crystal structure & properties

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previously

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Gases, Liquids, and Solids

Easily compressedExpand spontaneouslyto fill container

Retain volumeConform to shape of the containerAble to flowNearly incompressible

Retain volumeRetain shapeVirtually incompressibleOften have crystallineshape

Widely spaced moleculeswith much empty spacebetween themRandom motionVery weak attractionsbetween the molecules

Molecules tightly packedbut little orderAble to move past eachother with little difficultyIntermolecular attractiveforces relatively strong

Molecules tightly packedand highly orderedMolecules locked inplaceVery strong molecularforces

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There are 3 physical state of a material

Inter vs Intra-Molecular Forces

Cl H Cl H

Covalent Bond (strong)

Intermolecular attraction (weak)

Intermolecular forces:The attractions between molecules

Intramolecular forces:The chemical bonds that hold molecules together

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The closer two molecules are, the more strongly they attract each other.

Van der Waals forces • Van der Waals forces driven by induced electrical interactions

between two or more atoms or molecules that are very close to each other.

• Van der Waals interaction is the weakest of all intermolecular attractions between molecules.

• However, if a lot of Van der Waals forces interacting between two objects, the interaction can be very strong.

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Types of Intermolecular ForcesVan der Waals forces

1. London Forces

2. Dipole-Dipole Attractions

3. Hydrogen Bonds

4. Ion-Dipole Attraction

5. Ion-Induced Attraction

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1. London forces

• Interaction between nonpolar substances

• Movement of electrons in one atom influences the movement of electrons in other atom.

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Movement of electron density can be unsymmetrical, cause more negative charge on one side that on the other DIPOLE or Instantaneous dipole

Induce unsymmetrical electron density of its neighbor induced dipoleCalled as instantaneous dipole-induced dipole interactions

2. Dipole-Dipole Interactions

• Polar molecules, which has permanent dipole.

• Called as dipole-dipole interactions

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• Weaker than covalent bonds, 1-4% as strong

• Fall off rapidly with distance

3. Hydrogen bonds

• Hydrogen covalently bonded to a very small, highly electronegative atom (F, O, or N)

• Strong dipole-dipole attraction

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4. Ion-Dipole Attractions

• Ion-dipole attractions: when ions attract the charged ends of polar molecules

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5. Ion-induced Dipole Attractions

• Ion can induce dipoles in neighboring particles ion-induced dipole attractions

• Stronger than London forces

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Ion-induced dipole forces

Intermolecular Forcesand Properties of Liquids and Solids

Properties that Depend on Tightness of Packing:• Compressibility

– a measure of ability of a substrate to be forced into a smaller volume

• Diffusion– Ability of molecules to move from one position to another– occurs more rapidly in gases than in liquids, and hardly in

solids

In liquids, a given molecule suffersmany collisions as it moves about, (takes longer to move from place to place, making diffusion much slower).

Diffusion in solids is almost nonexistent at RT because the particles are held tightly in place.Gases compress easily molecules are far apart.

Liquids are incompressiblemolecules arepacked together.

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Properties that Depend on Strengths of Intermolecular Attractions:

1. Retention of Volume and Shape

In liquids and solids, the attractions of molecules are stronger than gas hold the particles closely together:

As a result, liquids and solids keep the same volume regardless of the size of their container

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2. Surface Tension• The property of the surface of a liquid that allows it to resist

an external force, due to the cohesive nature of its molecules

• Molecule at the surface higher potential energy than a molecule in the bulk

• A system is more stable when its potential energy decreases.

• For a liquid, reducing its surface area lowers its potential energy.

• The smallest surface area possible is spherical shape.

http://socratic.org/questions/what-is-surface-tension

Like invisible “skin”

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Water droplet lying on a damask. Surface tension is high enough to prevent floating below the textile

3. Wetting of a Surface by a Liquid

• Wetting: the spreading of a liquid across a surface to form a thin film

• For wetting to occur: The intermolecular attractions betweenthe liquid and the surface = the attractions within the liquid itself.

• EX: when water touches clean glass. The glass surface contains lots of oxygen atoms H2O can form H-bonds

• When the glass is coated by a film of oil or grease interact by London forces very weak compare to H bonds.

The attractions within liquid water are much stronger than the attractions between water molecules and the greasy surface.

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4. Viscosity • Resistance to a change in form of liquid

• EX: syrup is more viscous than water flows less readily than water

• Viscosity is influenced by:

• Intermolecular attractions,

2. Molecular shape and size

3. Temperature.

Intermolecular attractions

• INCREASE of the strengths of the intermolecular attractions INCREASE the viscosity.

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dipole–dipole attractions + London forces

London forces + strong hydrogen bonding

polar carbonyl group

More viscous than acetone.

Molecular shape and size• The long, floppy, entangling molecules in

heavy machine oil (almost entirely a mixture of long chain, nonpolar hydrocarbons) has 600 times viscosity that of water.

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Temperature

• As the temperature DECREASES, the viscosity INCREASES

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Changes of State and Dynamic Equilibria

• Occurs when a substance is transformed from one physical state to another

Changes of state between gases, solids, and liquids

Toward a condition of dynamic equilibrium

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Changes of State Involve Equilibria(Equilibria in Evaporation)

Before system reaches equilibrium

• Liquid is placed in empty container begins to evaporate

• Once in gas phase, molecules collide with each other, walls of the container, and the surface of the liquid scattering their kinetic energies condensed.

• Before equilibrium, Rate of Evaporation >> Rate of Condensation

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System At Equilibrium

• Rate of evaporation = rate of condensation

• The number of molecules in the vapor will remain constant = dynamic equilibrium

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Equilibria in Melting• Melting Point : The temperature at

which the solid begins to change into liquid as heat added

• At this temperature dynamic equilibrium exists between solid and liquid states

• Molecules leave the solid and enter the liquid at the same rateas molecules leave the liquid and join the solid (freeze)

• As long as no heat is added or removed from such a solid–liquid equilibrium mixture, melting and freezing occur at equal rates

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Equilibria in Sublimation

• At equilibrium, molecules evaporate from solid at the same rate as molecules condense from vapor

Equilibrium is established when molecules sublime from the solid at the same rate as they deposit on the solid from the vapor.

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Evaporation

• One of the most important physical properties of liquids and solids is their tendency to undergo a change of state from liquid to gas or from solid to gas

• For liquids the change is called EVAPORATION• For solids, the change directly to the gaseous state by

evaporation without going through the liquid state is called SUBLIMATION .

• EX : Solid carbon dioxide “dry ice”• How? at a given temperature, molecules have very large

kinetic energies very high velocities escape the attractions of its neighbors and enter the vapor state.

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Cooling Effect by Evaporation

when molecules evaporate they carry with them large amounts of

kinetic energy. As a result, the average kinetic energy of the molecules left behind decreases reduce the temperature of remaining liquid

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Rate of EvaporationRate of evaporation depends on factors :

• surface area of the liquid: when the surface area is increased, more molecules are able to escape liquid evaporates more quickly

• Temperature: As the temperature increase, Kinetic energy of the

molecules increases, and the rate of evaporation is also increase

• Strengths of intermolecular attractions: the stronger the intermolecular attractive forces, the slower is the rate of evaporation at a given temperature.

.

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Vapor Pressures of Liquids and Solids

• At equilibrium, the concentration of molecules in the vapor remains constant and the vapor exerts a constant pressurecalled as the equilibrium vapor pressure

http://dluetgens.com/vapor_pressure_equilibrium.html

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• Vapor pressure : the pressure caused by molecules that enter the vapor when a liquid evaporates

Factors affecting the equilibrium vapor pressure

1. Temperature:

The temperature increases The vapor pressure increases.

2. The intermolecular forces

The stronger the intermolecular forces The lower the vapor pressure.

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Factors that are not affecting vapor pressure

• Vapor pressure doesn’tdepend on the total SURFACE AREA, or on the VOLUME of liquid, or on the VOLUME of container.

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• Those factors don’t affect the rate of evaporation per unit surface area.

Boiling Point

• Boiling point depends on atmospheric pressure 1 atm chosen as the reference pressure

• strong intermolecular attractions causing low vapor pressure results in High boiling point

A liquid at its boiling point

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• The boiling point is defined as the temperature at which

the vapor pressure of the liquid = atmospheric pressure.

Energy & Changes of StateHeating Curve

In the heating curve, changes of state will increase the distance between molecules increase the potential energy.

Increasing temperature increase the kinetic energies of the particles.

If temperature is constant, there is no changes in kinetic energy. All of the heat will increase the potential energies.

Superheating/supercooling : occurs when the liquid is heated/ cooled above the boiling point/below the freezing point.

Cooling Curve

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• The potential energy changes associated with melting and vaporization expressed as enthalpy changes ∆H (molar heat per mole):

1. molar heat of fusion, ∆H fusion: the heat absorbed by one mole of a solid when it melts to give a liquid at the constant temperature and pressure.

2. the molar heat of vaporization, ∆𝐻 vaporization: the heat absorbed when one mole of a liquid is changed to one mole of vapor at a constant temperature and pressure.

3. the molar heat of sublimation, ∆𝐻 sublimation: the heat absorbed by one mole of a solid when it sublimes to give one mole of vapor at a constant temperature and pressure.

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How to calculate the heat required to melt specific amount of a substance?

q = heat required and n is the number of moles.∆H sublimation; ∆H vaporization; ∆H fusion

http://wps.prenhall.com/wps/media/objects/4678/4790699/ch08_07.htm

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q= heat required (J)m= mass (g)s= specific heat capacity (J.g-1.0C-1)Δt= changes in temperature (0C)

1. heat the solid to the melting point

2. melt the solid into liquid

3. heat the resulting liquid until 515 °C

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1.

1

2

3

97.80C

5150C

250C

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2.

3.

4.

• Evaporation / sublimation cause reduction in attractive forces between molecules.

• ∆H = energy needed to separate molecules

• Large ∆H = strong intermolecular attractions

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Le Chatelier’s Principle

“When a dynamic equilibrium in a system is upset by a disturbance, the system responds in a

direction that tends to counteract the disturbance and, if possible, restore the equilibrium”

http://ths.talawanda.org/~bramblen/classroom/Chemistry/Notes/Section%205A/LeChateliersPrinciple.htm

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Adding heat will raise the temperature of the equilibrium system, thus the system will try to adjust in a way that absorbs some of the added heat.

When liquid evaporates , the amount of vapor increase and cause the pressure to raise.

Disturbance in a system changes the position of equilibrium “equilibrium has shifted”

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Phase Diagram

• Phase diagram :

– graphical representation of the pressure-temperature relationships that apply to the equilibria between the phases of the substance

Solid-vapor equilibrium line

Liquid-vapor line

Triple point

Phase diagram of Water

Melting point line

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• Critical temperature: the highest temperature at which it is possible to separate substances into two fluid phases (vapor and liquid): point D

• critical pressure: P at point D pressure required to keep the gas in the liquid form at critical temperature

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Supercritical Fluids

A substance that has a temperature and pressure aboveits critical temperature and pressure and have density near its liquid density is described as a supercritical fluid.

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Low density

High densityDensity decreased

Density increased

Density decreased

Density increased Density is identical

SUPERCRITICALFLUID

• No distinct phase exist above the critical T and P.

• Supercritical fluid has density HIGHER than average gas, and LOWER than average liquid

• supercritical fluid have properties of both gas and liquid states able to dissolve like liquid, and effuse like gas

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https://www.youtube.com/watch?v=GEr3NxsPTOA

supercritical fluids

Structures of Crystalline Solids

• As substances freeze, or when they separate as a solid from a solution, They form crystals with highly regular features high degree of order among the particles that lie within the crystal

• The symmetrical features of a repeating structure, called lattice

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The simplest and most symmetrical three-dimensional lattice: simple cubic

Lattice = Unit cell: small repeating entity of the atomic structure. The basic building block of the crystal structure.

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The Structure of solid crystals

• Simple cubic

• Body-centered cubic

• Hexagonal closest-packed

• Cubic closest-packed

http://chemwiki.ucdavis.edu/Physical_Chemistry/Physical_Properties_of_Matter/Phases_of_Matter/Solids/Crystal_Lattice/Closest_Pack_Structures

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Crystal Types and Physical Properties

• Ionic Crystals

• Molecular Crystals

• Covalent Crystals

• Metallic Crystals

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Ionic Crystals

• Have ions at the lattice sites, the binding between them is mainly electrostatic.

• Tend to be hard

• Have high melting point. Need a lot of kinetic energy to enable them entering liquid state.

http://kraftyarts.com/crystals-beauty-art-and-healing/

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Ex: NaCl, KCl, KF, KI

Molecular Crystals

• The lattice sites are occupied either by atomsor by molecules.

• Soft

• Have low melting points

• Example : solid argon, krypton, H2O, SO2

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Covalent Crystal• Lattice positions are occupied by atoms that are

covalently bonded to other atoms at neighboring lattice sites.

• Tend to be very hard

• Have a very high melting point

The structure of diamond, an example of a compound that consist of covalent crystals.

http://dreamatico.com/diamond.htmlhttps://en.wikipedia.org/wiki/Diamond

Graphite and Diamond Crystal Structures

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Metallic Crystals

• Conduct heat and electricity

• Melting points range from high to low

• Range from very hard to very soft

The “electron sea” model of metallic crystal

Stibnite (Sb2S3) with barite (BaSO4)

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Summary

• Types of intermolecular forces :– london forces, dipole-dipole attraction, hydrogen

bonds, ion-dipole attraction, ion-induced attraction.

• Properties depend on tightness of packing– Compressibility, diffussion

• Properties depend on intermolecular force– Retention of volume and shape, surface tension,

viscosity, wetting of a surface by liquid

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Summary

• Changes of state always involve equilibria

• Intermolecular attractions also affect energy changes

• Le Chatelier’s Principle

– When an equilibrium is disturbed, the system will adjust the condition to counteract the disturbance

• Crystal Types

– Molecular, ionic, covalent, metallic

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Reference

• Jespersen, N.D., J. E. Brady, & A. Hyslop. 2012. Chemistry: The Molecular Nature of Matter 6th

Edition (Chapter 12)

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