Chapter 12 Intermolecular Forces Earth: 15 ºC Uranus: -214 ºC

Chapter 12Intermolecular Forces

Earth: 15 ºC

Uranus: -214 ºC

What happens to molecules at the melting point?

Intramolecular forces versus Intermolecular forces (aka. van der Waals forces)

Chemical Change

Breaks covalent, ionicand metallic bonds

Physical Change

Electrostatic forcesbetween particles

Intramolecular forces

Intermolecular forces

Section 12.1: Physical States and Phase ChangesKinetic-Molecular View of Three States of Matter

Increasing Energy

Solid Liquid Gas

P.E. and K.E. together determine the physical state

of any given substance.

P.E. – draws molecules together

Coulomb’s Law (Chap2) – particles withopposite charge attract each other.

The energy of attraction between twoparticles is proportional to the product of

the charges and inversely proportionalto the distance between them.

K.E. – separates or disperses molecules

K.E. ~ f(absolute temperature) (in Kelvins)

ºC = K - 273

Increasing Energy

Solid Liquid Gas

• Definite shape.• Definite volume.• Particles fixed & close.• Particle interaction v. strong.• Particle movement v. slow.

Ex: ice, iron, table salt

• Takes shape of container.• Definite volume.• Particles (molecules) random & close.• Particle interaction strong.• Particle movement moderate.

Ex: water, oil, vinegar

• Takes shape of container.• Fills container volume.• Particles (molecules) random and far apart.• Essentially no interaction. • Particle movement very fast. Examples: water vapor, helium gas

Section 12.1: Characteristics of Physical States

Section 12.1: Phase Changes

Condensation: Gas to liquidVaporization: Liquid to gas

Freezing: Liquid to solidMelting (Fusion): Solid to liquid

Sublimation: Solid to gasDeposition: Gas to solid

Heat of vaporization

Heat of fusion

Heat of sublimation

Section 12.1: Energy and Phase Changes

Enthalpy changes (∆H) accompany phase changes

Exothermic Phase Changes (-∆H) Condensation: Gas to liquid Freezing: Liquid to solid Deposition: Gas to solid

Enothermic Phase Changes (+∆H) Vaporization: Liquid to gas Melting (Fusion): Solid to liquid Sublimation: Solid to gas

Section 12.1: Energy and Phase ChangesEnthalpy change is different for different substances

For a pure substance: ∆H is measured in change per mole of the substance and is specific to the pressure and temperature conditions

Pressure is usually 1 atm, Temperature is that at which the phase change occurs

Example: Phase changes of water

H2O (l) H2O (g) ∆H = ∆H º vap = 40.7 kJ/mol (at 100 ºC)H2O (s) H2O (l) ∆H = ∆H º fus = 6.02 kJ/mol (at 0 ºC)

H2O (g) H2O (l) ∆H = ∆H º vap = -40.7 kJ/mol (at 100 ºC) H2O (l) H2O (s) ∆H = ∆H º fus = -6.02 kJ/mol (at 0 ºC)

Why is ∆H º vap (40.7 kJ/mol) greater than ∆H º fus (6.02 kJ/mol)?

∆H º subl = ∆H º fus + ∆H º fus

Section 12.2: Quantifying Phase ChangesT

emp

erat

ure

Heat Removed

The Heating-Cooling Curve – shows how the temperature of a substance changes asheat is added or removed from a substance at a constant rate (at a constant P too)

Interlude: Pressure Matters too (but we assume 1 atm in this class for phase change calculations)

http://pathways.fsu.edu/faculty/geeo/

Interlude: Pressure Matters too (but we assume 1 atm in this class for phase change calculations)

Methane hydrates

http://www.windows.ucar.edu/tour/link=/earth/Water/temp.html&edu=mid

Pressures ~ 1000 atmCH4 freezing point: -182.5 ºC

Section 12.2: Quantifying Phase ChangesT

emp

erat

ure

Heat Removed

But back to temperature……….. 5 heat-releasing stages

Temp change: q = nC∆T where q is heat, n is # of moles, C is molar heat capacity

Temp constant: q = n∆H where q is heat, n is # of moles, ∆H is heat released/absorbed

Section 12.2: Quantifying Phase Changes

In class problem: 12.20

Suggested problem: 12.2. 12.3, 12.12, 12.19, 12.27

Section 12.2: Equilibrium and Phase Changes

In a closed system, phases changes of many substances reach equilibrium.

Open Container Closed Container

Open system – volume of liquid does not change – net direction of molecule movement is out of the liquid

Closed system – volume of liquid does not change – net direction of molecule movement is out of the liquid

Closed Container

Systems can be closed tosome things, but not others

Heat source

Your system is defined by you:

Is the ocean a closed system?

Is Earth a closed system?

When matter/energy is leaving and entering an open system, it can reach Steady-state

Water in

Water out

Definition of steady-state: FluxIN = FluxOUT

Flux: Mass or Volume / time

Ex: 100 L H2O / hr 20 g CaCO3 / day

Water in

Water out

Water in

Water out

Water in

Water out

Time

This system is not in steady-state if the volume changes with time.

Concept of Open Systems and Steady-State

Vapor is stuck in the container and will accumulate, putting pressure (P=Force/Area) on container walls.

Section 12.2: Liquid-Gas Equilibrium and Vapor Pressure

Equilibrium vapor pressure – the pressure exerted by a vapor when it has reachedequilibrium in a system that is closed with respect to the vapor molecules

When enough time passes,the system will reach equilibriumwith respect to the vapor enteringand exiting the liquid.

Universal Concept: When a system at equilibrium is disturbed, it counteracts thedisturbance and eventually re-establishes a state of equilibrium (For chemical reactions, called Le Châtelier’s principle Chap17, CHEM 163)


Higher T = Higher V.P.Higher T increases the fraction of molecules moving fast enough to escape the liquid

decreases the fraction of molecules moving slow enough to be captured


Clausius-Clapeyron equation - mathematical relationship between T and P

Nonlinear relationship between T and P(in graph) expressed as linear relationship:

Know: P1,T1, ∆Hvap

Section 12.2: Quantifying T – P Relationships

In class problem: 12.22, 12.24

Suggested problem: 12.21, 12.23

Pressure exerted on the Earth by all the gas particles in the Earth’s atmosphere.

Pressure = Force / Area

From physics: F = ma F force m mass of particle a acceleration (= g, acceleration due to gravity)

Atmospheric Pressure

Section 12.2: Vapor Pressure and Boiling Point

Boiling point – temperature at which the vapor pressure equals the external pressure

Altitude = 10,000 feetAtmospheric pressure lower = 590 mm Hg

Boiling Point = 90 °C

Altitude = Sea level760 mm Hg (=1 atm)

Boiling Point = 100 °C

People living in Denver, CO use pressure cookers to cook food at higher temperature.

Why is my soup not as hot at Camp Muir?!??!

Section 12.3: Types of Intermolecular Forces

Bonding (Intramolecular) forces: Relatively strong involve large charges that are closer together

Nonbonding (Intermolecular) forces: Relatively weak involve smaller charges that are farther together


Why are bonding (intramolecular) forces strongerthan van der Waals (intermolecular forces)?

Periodic Table trends are similar tothose for bond length.


(1) Ion-dipole – an ion interacts with a partial charge

Dissolution(NaCl “dissociates”)

++

+ +

++

++

Na+1

++

+ ++

+

++

Cl-1

Na Cl Na Cl

Na ClNa Cl

NaClNaCl

NaCl NaCl

Example: NaCl (table salt) dissolves in water

H H

O

+ +


(2) Dipole-dipole – polar molecules interact

The greater the dipole moment of a molecule, the great the dipole-dipole forces between molecules of that type more energy needed to separate them


(2) Dipole-dipole – polar molecules interact

Hydrogen bond – a special type of dipole-dipole force that arises between atoms thathave a H atom bonded to a small, highly electronegative atom with lone electron pairs

N, O, and F all fit this profile.

H H

O

+ +

(3) Charge-induced dipoles – a molecule with a full or partial charge induces atemporary dipole on a nonpolar molecule

++

++

++

++

(4) London (dispersion) forces – caused by momentary oscillations of e- chargein atoms and, therefore, are present in all particles (atoms, ions, and molecules)

Section 12.3: Trends in Polarizability

Polarizability – the ease with which the e- cloud of a particle can be distorted

Smaller atoms (ions) are lesspolarizable than larger ones e-’scloser to the nucleus and, therefore,held more tightly

Polarizability• Increases down a Group

• Decreases from L R

• Cations less polarizable than their original atoms

Anions are more polarizable than original atoms

Why dry ice (solidCO2) sublimates

Biodiesel Lab: Bomb Calorimeter

Combustion reaction – heat flows from the system to the surroundings = exothermic

Heat is lost to: (1) water in the calorimeter (2) the calorimeter itself

Increasing Energy

Solid Liquid Gas

Section 12.4: Zooming in on Liquids

Randomness of particles any region is pretty muchidentical to any other

Orderliness of particles Different regions identical

Liquids are least understood at the molecular level.

Orderly &random at

different times

Macroscopic properties of liquids are well understood:• Surface tension• Capillarity• Viscosoty

• At the surface of a liquid, water molecules behave as a thin, elastic membrane or “skin” surface tension – energy required to increase the surface area (J/m2 of surface area increased)

How insects walk on water. (water strider)

Section 12.4: Surface Tension

Intermolecular forces exert different effects on a molecule at the surface of a liquidthan at the interior: A liquid tends to minimize the # of molecules at the surface.

Interior molecules – attracted by watermolecules on all sides

Surface molecules –attracted to water molecules below and on sides Experience a net downward attraction

Surfactants (surface- active agents)—destroy surface tension by congregating at thesurface and disrupting the hydrogen bonds between water molecules

Example: Needle on water.

H-bonds break,needle sinks.

H-bonds hold needleon water surface.

Example: Respiratory distress syndrome (RDS) in infants.Occurs when infantdoes not produce asurfactant that breaksH-bonds and does not allow O2 and CO2

exchange betweenalveoli and capillaries inthe lungs.

Section 12.4: Surface Tension

The stronger the forces arebetween the particles in a liquid,the greater the surface tension.

Section 12.4: Capillarity

Capillary action – the rising of a liquid through a narrow space against the pull ofgravity due to competition between intermolecular forces in a liquid (cohesive forces)And those between the liquid and the tube walls (adhesive forces)

TLC and plant pigment lab

Meniscus on a test tube

Glass = SiO2

Water(H-bonds with SiO2)

Mercury(Metallic bonds stronger than

any interaction with SiO2)

Section 12.4: Viscosity

A liquid’s resistance to flow resistance decreases as Temp increases

Molecular shape plays a role – Biodiesel lab

triglyceride + methanol 3 methyl ester + glycerol

Smaller molecules – make less contact = lower viscosity

Larger molecules – make more contact = higher viscosity

Section 12.5: Uniqueness of Water

H H

O

+ +

• •

The water molecules is bent and highly polar due to this structure and charge distribution, water

can engage in four H bonds with its neighbors.

(1) Water is the “universal solvent” (solvent = the compound that does the dissolving)

Dissolves a range of solutes ( = the compounds that are dissolved)

Na Cl Na Cl

Na ClNa Cl

NaClNaCl

NaCl NaCl

Ionic substances Polar Covalentsubstances

CH3CH2OHC6H12O6

Nonpolar Covalentsubstances

N2 gas


(2) Water has a high specific heat capacity (the measure of the heat absorbed by a substance for a given rise in temperature – Section 6.3)

In other words, water can absorb a lot of heat with relatively small changes in temp.

Earth: Daily temperature changes = 40 ºC (in deserts – most extreme)Waterless Moon: 250 ºC daily fluctuations

Water has a high heat of vaporization – heat from Sun results in vaporization ofocean water heat stored in water vaporcarried poleward heat released when water vapor condenses back to liquid water – called latent heat transport

oceanmotion.org/html/background/climate.htm

http://oceanmotion.org/html/background/climate.htm


(3) Surface properties are crucial to living things

Solids(minerals)

Air

Water

Trees get water due to capillary actionin soils and in xylem (veins of trees)

Plant debris floating on water surfaceprovides shelter and nutrients


(4) Density of solid and liquid water

Large spaces in the ice due to the hexagonal crystal structure result in solid waterbeing more dense than liquid lake surfaces freeze in winter (organisms live below)

Section 12.5: Uniqueness of Water Summarized

Heating-Cooling Curve Practice

How much heat would need to be added to heat 50.0 g of water ice at -50.0 ºC towater vapor at 135 ºC?

Given:

Cice = 37.6 J/mol ºCCliquid = 75.4J/mol ºCCgas = 33.1 j/mol ºC

∆Hfusion = 6.02 kJ/mol ∆Hvaporization = 40.7 kJ/mol

Answer: 1.59 x 105 J

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Chapter 12 Intermolecular Forces Earth: 15 ºC Uranus: -214 ºC