Chapter 11Intermolecular Forces,

Liquids, and Solids

John D. Bookstaver

St. Charles Community College

St. Peters, MO

2006, Prentice Hall, Inc.

Chemistry, The Central Science, 10th editionTheodore L. Brown; H. Eugene LeMay, Jr.;

and Bruce E. Bursten

11.1

Molecular Comparison of

States of Matter

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.

States of Matter & Kinetic Energy

• The state of matter of a substance at a particular temperature and pressure depends on two quantities:Kinetic energy of particles

• KE is directly proportional to temperature• At higher temps, more KE = faster moving

particles = weaker interactions/attractive forces between particles

Strength of the attractions between particles• Forces of attraction bond particles together• Stronger the attractions = closer together

particles are to each other

The States of Matter

Both of these quantities (KE & attractive forces) are present in every substance, in every phase

• In gases, the KE is great enough to overcome forces of attraction that keep the molecules bound together (thus they diffuse or spread apart from one another)

The States of Matter• In liquids, less KE means the forces of

attraction are greater than in gases, thus liquid particles are closer together, move slower but are still able to flow past one another

• In solids, there is very little KE (only vibrational energy) and the attractive forces are the greatest within a solid, which keeps the particles bonded together in ridged crystalline structures

Phase Changes: Two ways to change state of matter

• Change in temperature:Increasing the thermal energy of a substance

allows it to have a greater kinetic energy (on a molecular level)

This means, if heated enough, the attractive forces which hold the particles of solid & liquids together can be overcome by KE

Melting: solid liquidVaporization: liquid gas

Phase Change: Two ways to change state of matter

• Change in pressure: This only affects gases because they are

compressibleApplying pressure will force the particles closer

to one another, increasing the strength of attraction between particles

Condensation: gas liquid

11.2

Intermolecular Forces (IMFs)

• Why are some substances gases at room temperature, but others are liquid or solid?

…it all goes back to chemical bonding and the strength of forces of attraction (strongest bonds form solids at room temp)

• 2 categories of bonding/attractive forces:

INTRA and INTER

INTRAmolecular Bonding

• Extremely strong bonding between atoms within a molecule (responsible for holding compounds together)Ex) O-H bonds within a water molecule, Na-Cl

bonds within sodium chloride.Ionic (M & NM) opposite charges attracted to

one anotherCovalent (NM & NM), and Metallic (M & M)

sharing of e– (nucleus of one atom is attracted to e– of another atom)

INTERmolecular Forces (IMFs)• Attractions between molecules

Ex) how an H2O interacts w/another H2O molecule

• IMFs are always WEAKER than intramolecular bonds

INTERmolecular Forces (IMFs)

• The strength of a substance’s IMFs greatly impact many physical properties including boiling & melting point, vapor pressure, & viscosity

The stronger the IMFs, the more energy is required to overcome the attraction between the molecules (this means higher melting & boiling points)

4 Types of IMFs

• There are 4 types of IMFs One type of IMF involves charged ions (ion-dipole forces) . The other three IMFs involve neutral molecules (dipole-dipole forces, London Dispersion Forces, & hydrogen bonding - these three are collectively referred to as the van der Waals forces).

Ion-Dipole Forces (IDFs)• A 1st type of IMF, ion-dipole interactions, involve

charged molecules are an important force in forming aqueous solutionsAttractions between ions and polar molecules

(polar = partial charges)

Example) NaCl + H2O

Since polar molecules are asymmetrical, they have partial charges where one side is partially positive (δ +) and one side is partially negative (δ–).

Ion-Dipole Forces (IDFs)

Cations are surrounded by the negative end of polar molecules (vice-versa for anions)

IDFs explain how

water (or other polar solvents) can interact with and dissolve soluble ionic compounds

Dipole-Dipole Forces (DDFs)

Dipole (2 poles) is another way to describe an atom or molecule which has a separation of charges (half positive, half negative)

The positive end of one polar molecule is attracted to the negative end of another molecule near by (and vice versa)

• Attractions between neutral, polar molecules

Example) PCl3 + PCl3

Dipole-Dipole Forces (DDFs)

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

There are also repulsive forces that occur when like partial charges come close to one anotherExample) PCl3 + PCl3

Dipole-Dipole Forces (DDFs)

• For molecules of approx. the same mass & size, the more polar the molecule (higher dipole moment), the stronger the DDFs and the higher the boiling point.

London Dispersion Forces (LDFs)

• Attractions between nonpolar molecules Ex) He + He; H2 + H2

Attractions must exist between them because nonpolar gases can be liquefied (thus attracted, somehow, to each other)

• In nonpolar atoms and molecules are very symmetrical by nature, therefore no dipole (separation of charge) typically exists A dipole can, however be induced…

http://besocratic.colorado.edu/CLUE-Chemistry/activities/LondonDispersionForce/1.2-interactions-0.html

#2:

A BRIEF polarization of anatom or molecule occurs when e– TEMPORARILY (by a freak

occurrence) cluster on one side of the molecule

#1: Symmetrical Molecule

#2: Instantaneous Dipole

#3: Induced Dipole in neighboring atom

#1:

In symmetrical atoms or molecules, electrons are evenly dispersed

around the nucleus

#3:

The instantaneous polarization of one

atom causes the e– in another neighboring

atom to shift

London Dispersion Forces (LDFs)

#2) While the electrons in an atom would repel each other (and, therefore, tend to stay far away from each other), it does happen that the electrons occasionally wind up on the same side of the atom (by chance, for an instant)

London Dispersion Forces (LDFs)

#2) At that BRIEF instant, the helium atom is polarized, with an excess of electrons on one side and a shortage on the other side = instantaneous dipole (partial charges)

London Dispersion Forces (LDFs)

#3) As a neighboring atom comes nearby the “instantaneously polar atom” a dipole would then momentarily be induced in the neighboring atom…

This cycle continues with another neighbor…

London Dispersion Forces

Sum it Up:

London dispersion forces (LDFs) are attractions between an instantaneous dipole (which happens by chance) and an induced dipole (which happens when another nonpolar molecule is near an instantaneous dipole .

London Dispersion Forces

• **LDFs are present in all molecules, whether they are polar or nonpolar**

• The strength of LDFs depends on how easily e– in a molecule can be distorted (the moelcules’s polarizability).

Factors Affecting LDF Strength

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

• The larger the molecule (with a greater molecular weight), the more electrons it has, the stronger the LDFs, and the higher the boiling point of the substance

Hydrogen Bonding

• A special type of dipole-dipole force that occurs between polar molecules which include Hydrogen bonded to N, O, or F

• H-bonds are the strongest of the van der waals forces

Hydrogen Bonding• H-bonding arises

due to the high electronegativity of N, O, and F

•  When hydrogen is bonded to one of these very EN elements in a molecule, the more EN element pulls the electron further away from H, creating partial charges within the molecule (H is partially positive); this creates a very large EN difference within the molecule make it extremely polar

List the 4 IMFs in order of increasing strength

• #1) Ion-Dipole Forces - strongest• #2) Hydrogen Bonding• #3) Dipole-Dipole Forces• #4) London Dispersion Forces - weakest

• How does the strength of the IMFs impact properties like the boiling point of the substance?

The stronger the IMF, the higher the BP

•  How does the size (molecular weight) of an atom or molecule impact the strength of IMFs?

The larger the atom or molecule, the stronger the IMFs

How Do We Explain This?• The nonpolar series

(SnH4 to CH4) follows the expected trend smaller mass = lower BP.

• The polar series (top) follows the trend from H2Te through H2S, but water is quite an anomaly. Why is the BP so high for water?

* Note: polar molecules tend to have higher BP than nonpolar molecules of similar size due to DDFs vs. LDFs

Graph of Compound Boiling Points vs. molar mass

Summarizing Intermolecular Forces

Increasing strength of forces

11.3

Some Properties of

Liquids

Intermolecular Forces Affect Many Physical Properties

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

Viscosity• Resistance of a liquid to flow is

called viscosity. Viscosity is related to how easily

molecules can move past one another other.

The more viscose the substance, the greater its resistance to flow

• Viscosity increases with stronger IMFs More attractions between molecules = less KE =

less ability to flow• Viscosity decreases with higher temperature

Increase in KE which overcomes IMFs in the liquid

Surface Tension

• Surface tension is a property of a liquid that allows it to resist the pressure of an external force

• The higher the surface tension, the greater the resistance to external forces, and thus the stronger the IMFs holding the liquid together

Surface Tension• Within a liquid, all the molecules

on the interior of the liquid are attracted equally in all directions

• However, molecules on the surface of a liquid can only be attracted to molecules below themThe net downward force of

attraction acting on surface molecules causes some molecules to be pulled into the interior of the liquid, thus the molecule leaves the surface

Surface Tension• As molecules leave the surface of the liquid, the

surface area that the liquid takes up is reduced • As surface area is reduced, those molecules which

are left at the surface are forced closer together, which increases the IMFs between themThus, the stronger the IMFs between molecules,

the less the surface area, and the greater the surface “tension” (or resistance to outside forces) of the liquid

• Because the primary IMFs between water molecules is hydrogen bonding (very strong), water has an extremely high surface tension

Cohesive Forces of Attraction• Cohesion:

A general type of intermolecular attraction between molecules of the same substance

H2O interaction w/H2O (or Hg atoms w/Hg atoms)

• Surface tension is an example of cohesion…This is because H2O molecules are attracted to the

other H2O molecules directly below them

Cohesive Forces: Interrupted• Surfactants, such a soaps, are

primarily nonpolar compounds• The have a small, polar

“hydrophillic” head and a very long nonpolar, “hydrophobic” tail (long chain of carbons and hydrogens)

• When surfactant (soap) molecules are introduced into water, the H2O molecules nearest to the surfactant are separated (their cohesive forces are interrupted due to the nonpolar nature of the surfactant)

How does soap “clean”?• This is the chemical reason why we

use soaps to clean our hands, dishes, clothing etc…

• Things like bacteria, dirt, and oil are all nonpolar substances Water, being polar, is unable to

interact with these substances (nonpolar substances cannot have any dipole-dipole interactions with water)

Soap molecules will surround the bacteria/dirt with its nonpolar tail, leaving the polar portion of the soap exposed. Water can then come wash away this surrounded bacteria/dirt because it can interact with the polar portion of the soap

“Like Dissolves

Like”

Adhesive Forces of Attraction• Adhesion:

A general type of intermolecular attraction between molecules of different substances with similar polarities (typically between a liquid & solid)

Think of adhesive tape or glue (sticks 2 different things together)

Because glass is polar and water is polar, the molecules of the substances have dipole-dipole interactions

• Capillary Action is an example of adhesion

Capillary Action

Water will rise up a thin glass tube (opposite direction of gravitational pull)

This is due to the adhesive forces between polar water and the polar glass molecules• the water molecules on the

sides are attracted to the glass and “creep” up the side, pulling additional water molecules along through cohesive forces of attraction

In capillary action, the adhesive forces must be greater than the cohesive forces to pull water molecules away from itself

http://www.youtube.com/watch?v=cbTwMoNOwag

http://www.youtube.com/watch?v=-

X0u_le2xxs - extra

Cohesive vs. Adhesive Forces

• Meniscus in H2O and Hg:

water has a concave meniscus due to greater adhesive forces that attract the water surface edges to the glass

Hg has a convex meniscus due to its greater cohesive forces (Hg is nonpolar and metallic while glass is polar covalent there are not adhesive forces between them)

11.4

Phase Changes

Phase ChangesFill in the blanks with the

appropriate phase change

Condensation

Deposition

Melting/Fusion

Vaporization/Evaporation

Sublimation

Freezing

Phase Changes

Exothermic (ΔH = –)

Changing to a phase w/ less KE … energy must be released

• Fusion (melting)• Vaporization

(evaporation)• Sublimation

• Freezing• Condensation

• Deposition

• Each phase change has an opposite processThe magnitude of ΔH is the SAME for opposing

processes, but the sign is the opposite!

Endothermic (ΔH = +)

Changing to a phase w/ more KE … energy must be absorbed

Energy Changes Associated with Changes of State

• Heat of Fusion (ΔHfus)Energy required to change a solid into a

liquid at its melting point (MP)• Heat of Vaporization (ΔHvap)

Energy required to change a liquid into a gas at its boiling point (BP)

• Heat of Sublimation (ΔHsub)Energy required to change a solid directly

into a gasSum of ΔHfus and ΔHvap

Energy Changes Associated with Changes of State

• ΔHvap values are higher than ΔHfus b/c when the molecules are transitioning from liquid gas, nearly all IMFs must be overcome (vs. solid liquid, many IMFs still remain in the liquid phase)

Heating Curve of WaterAB = heating ice

BC = fusion/ melting

CD = heating liquid water

DE = vaporization/ evaporation

EF = heating steam

5 Steps to change ice into steam:

Heating Curve of an Unknown (pgs. 4-5)

Heating Curves• The heat added to the

system at the melting and boiling points is used to overcome IMFs (pulling the molecules farther apart from each other = change in potential energy).

• As a result, the temperature of the substance does not rise during the PHASE CHANGE.Appear as plateaus on

the heating curve

Calculating ΔH for Temperature & Phase Changes

• Since AB, CD, and EF are just heating single phases the energy required (q) can be determined by

q = mCpΔT Note: each state of matter has a unique specific

heat• Since BC and DE (plateaus) are phase changes, the

temp remains constant The energy requires (q) can be determined by

q = (mol)ΔHfus or q = (mol)ΔHvap respectively

Note: Depending on units of ΔHfus/vap, mass may be needed instead of mols

11.5

Vapor Pressure

Vapor Pressure• Some liquid molecules gain enough KE to

overcome liquid IMFs these molecules can vaporize into the gas phaseNot all molecules have enough KE to escape

liquid phaseSome that do escape will be attracted to other

liquid molecules and re-condense

Vapor Pressure• What is vapor pressure?

The pressure exerted by a vapor in equilibrium with its liquid phase. This occurs when the rate of evaporation is equal to the rate of condensation (so that there is a constant # of molecules in the vapor phase, thus exerting a constant force on the walls of a container).

Vapor Pressure & Boiling• How does the normal boiling point (BP) of a

substance relate to its vapor pressure?A liquid boils when its vapor pressure equals the

external atmospheric pressureThe normal boiling point is the temperature at

which a substance’s vapor pressure is equal to the standard atmospheric pressure.

• What is the quantitative value of atmospheric pressure?1 atm (atmospheres)101.3 kPa (kilopascals)760 torr (equivalent to mm Hg)

Vapor Pressure Curves

Factors that AffectVapor Pressure

• At a given temperature, what is the relationship between IMF strength and vapor pressure?The stronger the IMFs, the lower the vapor

pressure – due to stronger attraction, fewer molecules can escape into the vapor phase

• What is the general trend of vapor pressure as temperature increases?As temp increases, vapor pressure increases –

because as temp rises, KE increases, meaning more molecules can escape into the gas phase by overcoming IMFs

Vapor Pressure• What does it mean if a

liquid is volatile?Volatile liquids

evaporate readily (have low BP) and, at a given temp, tend to have higher vapor pressures - because more molecules can easily escape to the vapor phase

Which substance in the diagram is the MOST volatile?

Vapor Pressure CurvesComparing Chloroform, Ethanol, Water, & Ethanoic Acid

11.6

Phase Diagrams

Phase Diagrams

• A phase diagram describes the temperature and pressure conditions at which a substance can exist as each state of matter.

• Each region represents a different state of matter.Solids

• exist at higher pressures & lower tempsGases

• exist at lower pressure & higher temps

Phase Curve Diagrams

• Each state of matter is separated by a line which gives the specific conditions when 2 phases exist in equilibrium with one anotherThese lines signifies when (under what

conditions pressure and temp) a phase change occurs

Phase Diagram

of Bromine

(Br2)solid

gas

liquid

MP BP

triple

“A”

“B”

“C”

Phase Diagrams

Phase Diagrams• The AD line is the liquid-solid interface (both

phases in equilibrium)• The melting point at each corresponding

pressure is found along this line.Normal MP = temperature at atmospheric

pressure

Phase Diagrams• The AB line is the liquid-vapor interface (both phases

in equilibrium).Each temperature along this liquid-vapor line is the

boiling point at the corresponding pressure.At 1.0 atm, the temperature is equal to the normal

boiling point

Phase Diagrams• At the end of the liquid-vapor line is the

critical point (B); above this critical temperature and critical pressure the liquid and vapor are indistinguishable from each other.

Phase Diagrams• The AC line is the solid-vapor interface (both

phases in equilibrium).Along the AC line the solid and gas phases

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

• Describe how the melting point of Br2 changes with external pressure.As pressure

increases, the MP increases. This is due to the increasing strength of IMFs as molecules are forced closer togetehr (higher MP needed to overcome stronger IMFs)

• Based on the slope of the melting-point equilibrium line and the fact that higher pressure (at the same temperature) always favors the more dense state of matter, would you characterize the solid phase of bromine as more or less dense than the liquid phase of bromine?For Br2, the MP equilibrium has a positive slope

(bends to the right)At a given temp (ex: 7°C)as pressures increases,

bromine would become a solid (thus, solid phase is more dense than liquid)

• What is the significance of the triple point?It describes the only set of conditions (for

temp & pressure)for which all 3 phases of a substance exist in equilibrium with each other

https://www.youtube.com/watch?v=r3zP9Rj7lnc

Phase Diagram of Water

• Triple Point of water:

low pressure4.58 torr =

0.006 atm

low temp 0.0098°C

Phase Diagram of Water

• The slope of the solid–liquid line of WATER 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