Che calc

Chemistry 1B - Foothill College

L. J. Larson- All rights reserved

Intermolecular Forces, Liquids and Solids1

Intermolecular Forces, Liquids,and Solids

• Understand intermolecular forces (IMF).

•Be able to determine the type(s) of IMF that exist for various substances.

•Understand properties of liquids and solids and how IMF effect theseproperties.

•Learn about the structure of solids and understand the differencebetween crystalline and amorphous solids.

•Understand the concepts of crystal lattice and cubic unit cells (LAB).

•Understand the different types of attractive forces in crystalline solids andhow these forces effect the properties of solids.


The States of Matter:Macroscopic Properties




States of Matter:Microscopic View

• A fundamental difference between states of matter is thedistance between particles.

• In the solid and liquid states particles are closer together, thuswe refer to them as condensed phases.


The strength of

attractive

forces between

particles

(atoms, ions or

molecules)

determines a

substance’s

phase at room

temperature.




The States of Matter

• The state a substance is in at a particulartemperature and pressure depends on twoantagonistic entities:

!The kinetic energy of the particles

!The strength of the attractions between the particles


Intermolecular Forces (IMF): Net Electrostatic AttractiveForce BETWEEN Molecules

• Directly related to properties such as melting point,boiling point and the energies to overcome theforces of attraction between particles in changes ofstate (!Hfus and !Hvap).

• Determines solubility of gases, liquids and solids invarious solvents.

• Crucial in determining the structure of biologicalmacromolecules such as DNA and proteins.

• IMF are short-range forces. They only exist overshort (nm) distances.




Energy and Intermolecular Forces

Energy is required to overcome intermolecular forces in a liquid

to make a gas. This energy is the heat of vaporization, !Hvap.


Intermolecular Forces

Intermolecular forces (IMF) as a group are generally referred to as van derWaals forces. These attractions are not nearly as strong as theintramolecular attractions that hold atoms together in compounds:

Let’s list approximate energies for the following:•Covalent Bond Energies:•Lattice Energies:•Intermolecular Force (IMF) Energies:




Intermolecular Forces

• IMF are, however, strong enough to control physical properties suchas:

• They are also responsible for deviations in ideal behavior of gases.


van der Waals Forces Include

• Dipole-dipole interactions

• Hydrogen bonding

• London dispersion forces




Dipole-Dipole Interactions (d-d)

• Molecules that havepermanent dipoles areattracted to each other.! The positive end of one is

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

! These forces are only importantwhen the molecules are close toeach other.

• Examples:


Dipole-Dipole Interactions

For molecules of about the same molecular weight, what can youconclude about the relationship between boiling point and molecularpolarity?

Can you explain this trend in terms of intermolecular forces?




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 happenthat they occasionally wind up on the same side of the atom.

• At that instant, then, the helium atom is polar, with an excess of electronson the left side and a shortage on the right side. An instantaneous dipoleresults.


London Dispersion Forces

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

• London dispersion forces, or dispersion forces, are attractionsbetween instantaneous dipoles.




London Dispersion Forces(induced dipole-induced dipole, id-id)

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

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

• Define Polarizability:


Factors Affecting London Forces

• From the above data, what can you conclude about therelationship between molecular weight and strength of LondonForces?

• Can you explain this relationship in terms of polarizability?




Factors Affecting London Forcesn-pentane neopentane

• What can you conclude about the relationship betweenmolecular shape and the strength of London Forces?Why do you think shape has this effect?


Br2 is a liquid at RT while I2 is a

solid. Why?

Br2 & I2




Dipole - Induced Dipole IMF (d-id)

A polar molecule will “induce” a TEMPORARY dipole in a

nonpolar molecule. The two molecules are now attracted

momentarily. This is typically a weaker interaction than d-d.

•Is this type of intermolecular force important to fish and other aquatic life?


Nonpolar gases are slightly soluble in water.

•Can you explain the trend in solubility shown above in terms of IMF?

•What would you predict for the solubility of He(g) compared to theabove gases?




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

• If two molecules are of comparable size and shape,differences in dipole-dipole interactions will likely bethe dominating force.Example:

• If one molecule is much larger than another,dispersion forces will likely determine differences intheir physical properties.Example:


How Do We Explain This?

• The nonpolar series(SnH4 to CH4)follows the expectedtrend: as molecularmass decreases,boiling pointdecreases.

• The polar seriesfollow the sametrend from periods 5to 3, but the period 2hydrocarbons areanomalies.

• Show H-bond video.

Why do these compounds

have such a high boiling

point?

Boiling points of some simple hydrogen compounds.




Hydrogen Bonding - A special case of d-d IMF.

• Here X is either a F, O or N atom and Y is a small highly electronegativeatom (N, O or F) that has a lone pair of electrons.

• When a molecule contains one of the following groups of atoms it can“hydrogen bond”:! O-H! N-H! F-H

• These groups are highly polar, the exposed H atom has almost a full +1charge. The resulting interaction of the H atom with a lone-pair of electronson a small, highly electronegative (N, O or F) adjacent atom gives anunusually strong d-d IMF. We call this a hydrogen bond.

The interaction can be described as:

X–H ----- :Y

hydrogen bond


HydrogenBonding

NOT a hydrogen bond

A hydrogen bond

Note: Hydrogen bonds areintermolecular forces.They are NOT covalentbonds!




Hydrogen bonding in ice.

•Does this explain why ice is less dense than liquid water?


Hydrogen bondingin acetic acidmolecules.

Notice how eachmolecule makestwo hydrogenbonds.




DNA & hydrogenbonding

phosphate-sugar

backbone

double helix

H-bonds


Ethanol & diethyl etherEthanol and diethyl ether are structural isomers. Predict which

substance has the greater boiling point.




Ion-Dipole Interactions

• A fourth type of force, ion-dipole interactions are an importantforce in solutions of ions.

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

•Examples:


Summarizing Intermolecular Forces




Like-Dissolves-Like!

Compounds with similar IMF are usually soluble in one another. This will be

covered in more detail in Chemistry 1C, Chapter 13.


Problem

• What IMF exist in the following systems?

! Pure CH4

! Water and CH3OH

! F2 and water

! HCl and CH3OCH3

• Rank the above systems in order of increasing strength of IMF




Intermolecular Forces Affect ManyPhysical Properties

The strength of the attractionsbetween particles can greatlyaffect the properties of asubstance or solution.

Liquids: Behavior is more difficult todescribe compared to gas behavior.

• Short-range order but no long-rangeorder

• Some properties of liquids:! Viscosity! Surface Tension

! Capillary Action! Heat of vaporization! Vapor Pressure

! Boiling Point! Critical Temperature and

Pressure


Viscosity

• Resistance of a liquid toflow is called viscosity.(SI units kg/m-s)

• It is related to the easewith which molecules canmove past each other.

• What does viscositydepend upon?




GlycerolWhich would you predict to have the greater viscosity, glycerol or ethanol,CH3CH2OH? Explain.

O

C

H


Surface Tension

• Surface tension results from the netinward force experienced by themolecules on the surface of a liquid.

• Surface tension is the energy needed toincrease the surface area of a liquid bya unit amount. Examples are:Water at 20°C 7.29x10–2 J/m2

Mercury at 20°C 4.6x10–1 J/m2

• What does surface tension dependupon?




Capillary Action: the ability of liquids to rise in tubes through

adhesive forces.

Adhesive forces:

Cohesive forces:


Phase Changes




Energy Changes Associated withChanges of State

• Heat of Fusion:

• Heat of vaporization:

• Using information from the above graph, what is the heat of sublimation ofbutane?

Stronger IMFgenerally resultin higher !Hfus

and !Hvap.


Energy Changes Associated withChanges of State

• The heat added to the system at themelting and boiling points goes intopulling the molecules farther apart fromeach other.

• The temperature of the substance doesnot rise during the phase change.

• What segments of the curve correspond toan increase in potential energy of thesystem?

• What segments of the curve correspond toan increase in kinetic energy of the system?

Heating Curve for water

Show change in state video.




•Text Problem 11.37: Drinking water can be cooled in hotclimates by evaporating it from the surfaces of canvas bags orporous clay pots. How many grams of water can be cooled from35°C to 22°C by the evaporation of 50.0 g of water. The heat ofvaporization of water in this temperature range is 2.4 kJ/g andthe specific heat capacity of water is 4.18 J/g-K.

•Which do you expect to have the greatest heat of vaporization?NH3 or PH3?

Cl2 or I2?


Evaporation and Temperature

• At any temperature, some molecules in a liquid haveenough energy to escape into the gas phase.

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




Vapor Pressure of a Liquid•As more molecules escape theliquid, the pressure they exert asa gas increases.

•The liquid and vapor reach astate of dynamic equilibrium:liquid molecules evaporate andvapor molecules condense at the

same rate.

!Define Vapor Pressure of a liquid:

!Vapor Pressure depends upon:

!Vapor Pressure does not depend upon:


Vapor Pressure of Butane

Show vapor pressure as a function oftemperature video




Boiling Point

• The boiling point of a liquid is the temperature at which its vapor pressureequals the external pressure exerted upon the liquid.

• The normal boiling point is the temperature at which a liquid’s vaporpressure is 760 torr (1 atm).

Estimate theboiling pointof water at:

600 mm Hg

1000 mm Hg


Problem:What is the vapor pressure of CS2 at 40 °C?

What is the normal boiling point of CS2?

What IMF are present for each compound?

Which compound has the strongest IMF?

Which liquid would have the lowest !Hvap?




Critical Temperature and Pressure

• Critical Temperature:

• Critical Pressure:

Which of thesesubstances can beliquefied at roomtemperature, about20°C?


Solids

• We can think of solids asfalling into two groups:(1) Crystalline(2) Amorphous

! Crystalline—particles(atoms, molecules orions) are in highly orderedthree dimensionalarrangement:IonicMetallicMolecularNetwork covalent

! Pure crystalline solidsexhibit “sharp” meltingpoints.




Solids

! Amorphous—no particularorder in the arrangement ofparticles.

! Amorphous solids do notexhibit “sharp” meltingpoints. Instead, theirtemperature changes asthey melt.


Crystalline Solids

•The particles (atoms, ions ormolecules) can be representedby a 3-dimensional array ofpoints called a crystal lattice.

•Because of the order in acrystal, we can focus on therepeating pattern ofarrangement called the unitcell.




Crystalline Solids: Overview of cubic unitcells covered in more detail in lab.

There are several types of basicarrangements in crystals, such as the onesshown above.


Types of Bonding in Crystalline Solids




Ionic Solids•Ionic forces of attraction

•Positive and negative ions

•Melting point tends to increase withincreasing lattice energy.Example: List in order of increasing mp;CaO, NaCl and KI

Properties

• Hard - brittle

• High melting points

• Poor electrical conductors assolid, good as liquid

• Many are water soluble. Aqueoussolutions conduct electricity.

Cl– ions primitive withCs+ ions in center

S2– ions fcc with the Zn2+ ions inhalf the tetrahedral “holes”

Ca2+ ions fcc with the F‒ ions in alltetrahedral “holes”


Ionic SolidsWhat are the empirical formulas for these compounds?

(a) Green: chlorine; Gray: cesium

(b) Yellow: sulfur; Gray: zinc

(c) Green: calcium; Gray: fluorine

CsCl ZnS CaF2

(a) (b) (c)




Covalent-Network SolidsAtoms held together in largenetworks of strong covalent bonds

Properties

• Hard

• High melting points

• Poor electrical and thermalconductors (usually)*

3.41Ådispersion forceshold layers together

mp = 3550°C

1.42Åcovalent bonds

*Graphite is a good conductor of electricity alongthe layers due to the delocalized " bond system


The unit cell for diamond is shown here.

What type of unit cell is this?




Covalent-Network Solids-Other Examples

Silicates - quartz, sand, etc.

SiO2 lattice


Molecular Solids

Forces of attraction are• Dispersion• Dipole-dipole• Hydrogen bonds

Properties

• Fairly soft

• Low to moderate melting points

• Poor electrical and thermalconductors

Ice -a molecular solid




Metallic Solids

• Metals are not covalently bonded, butthe attractions between atoms aretoo strong to be van der Waalsforces.

• In metals, valence electrons aredelocalized throughout the solid.

• Metallic forces of attraction (bonding):

“Electron Sea Model”Attraction between positive cores anda sea of valence electrons resultingin bonding that is non-directional innature.


Metallic Solids

Metallic Properties

• Malleable

• Ductile

• Good electrical and heatconductors

• Wide range of hardness

• Wide range of melting points,but many have high meltingpoints

Al is FCC




Amorphous Solids Properties

• Non crystalline

• Wide range of melting points

• Poor electrical conductors

Glass


Substance

Examples of Some Physical Propertiesof Solids

Metallic Bonding

Ionic Bonding




• Rank in order if increasing melting point: CH3F, MgO, H2O and CsI

• For each of the following pairs, predict which will have the higher meltingpoint and explain why.(a) HF or HCl (b) CH4 or CCl4 (c) KBr or H2O

• Given the following solids: BaCl2, Zn, CH3COOH and diamondWhich would conduct electricity?

Which would dissolve in water to give a solution that conducts electricity?

Which would be malleable?Which would be brittle?Which would have the lowest melting point?

Example Problems


Example Text Problems

11.7 Niobium (II) oxide crystallizes in the following cubic unit cell.

(a) How many niobium atoms and how many oxygen atoms are within theunit cell?

(b) What is the empirical formula of niobium oxide?

(c) Is this a molecular, covalent-network, or ionic solid?





11.16abd What type of intermolecular force accounts for the followingdifferences in each case?(a) CH3OH boils at 65°C, CH3SH boils at 6°C.

(b) Xe is liquid at atmospheric pressure and 120 K, whereas Ar is a gas.

(d) Acetone boils at 56°C, whereas 2-methylpropane boils at !12°C.



11.48 Explain the following observations:(a) Water evaporates more quickly on a hot, dry day than on a hot, humidday.

(b) It takes longer to cook hard-boiled eggs at high altitudes than at loweraltitudes.





11.102ab Liquid butane, C4H10, is stored in cylinders, to be used as a fuel.The normal boiling point of butane is listed as !0.5°C.(a) Suppose the tank is standing in the sun and reaches a temperature of46°C. Would you expect the pressure in the tank to be greater or less thanatmospheric pressure?

How does the pressure within the tank depend on how much liquid butaneis in it?

(b) Suppose the valve to the tank is opened and a few liters of butane areallowed to escape rapidly. What do you expect would happen to thetemperature of the remaining liquid butane in the tank? Explain.

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