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20190501234653ch11_lecture_notes1.pdf

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11.1 Kinetic Molecular Theory, States of Matter, and Phase Changes 11.2 Vapor Pressure 11.3 Other Properties of Liquids 11.4 The Nature of Intermolecular Forces 11.5 Intermolecular Forces and the Properties of Liquids

To this point, we have described the structures of individual atoms and individual molecules in great detail. For atoms, we have focused on electronic structure, and

for molecules, we have studied the arrangement of atoms with respect to one another. We are now ready to move to the next level of complexity in chemical

systems: how collections of molecules interact and how those interactions control the physical properties of matter.

Chapter 11. Intermolecular Forces and the Liquid State

Intermolecular Forces and the Liquid State

11.1 Kinetic Molecular Theory, States of Matter, and Phase Changes

• KMT also applies to liquids and solids

– Molecules in liquids and solids are in constant random motion

• Liquids and solids are called condensed phases

– Particles are packed in close proximity

– Density is a measure of the differences between the phases

– Except for water, the density increases when going from solid to liquid to gas

KMT for Solids and Liquids11.1

• Forces between particles that hold one molecule near another molecule

– Relative strengths of IMFs closely mirror that of the density ranking

– Greatest in solids and the weakest in gases

Intermolecular Forces (IMFs)

11.1

Physical

state

IMFs

between

particles

Compressibility Shape and

volume

Ability

to flow

Gas Generally weak High Takes on shape

and volume

of container

High

Liquid Generally

intermediate

Very low Takes on shape of

container;

volume limited by

surface area

Moderate

Solid Generally

strong

Almost none Maintains own

shape and

volume

Almost

none

Table 11.1.1 - Properties of Solids, Liquids, and Gases

Influence chemistry, directly related to chemical and physical properties: – Melting point

– Boiling point

– Energy to convert a solid to liquid

– Energy to convert liquid to vapor

– Solubility of gases, liquids, and solids in various solvents

– Structures of biologically important molecules, such as DNA and proteins

– Much weaker than covalent bonds

Intermolecular Forces11.1

Intermolecular Forces and the Liquid State

11.4 The Nature of Intermolecular Forces

11.4 Intermolecular Forces

Attractive forces that hold particles together in the condensed phases are called Intermolecular Forces (IMFs).

Magnitude of intermolecular forces is what determines if the particles in the substance are in a gas, liquid, or solid phase.

van der Waals forces -dipole-dipole interactions, including hydrogen bonding and dispersion forces

11.4 Intermolecular Forces

Dipole-dipole interactions are attractive forces that act between polar molecules.

As the dipole forces increase, intermolecular forces increase. As the intermolecular forces increase, boiling points increase.

11.4 Intermolecular Forces

Hydrogen bonding is a special type of dipole-dipole interaction between H-F, H-O, and H-N ONLY .

H bonded to a small, highly electronegative atom, such as N, O, or F .

• H-bonding is especially strong in water O—H bond is very polar

• 2 e- lone pairs on the O atom attract H-atoms from other H2O molecules

• Accounts for many of water’s unique properties.

Hydrogen Bonding in H2O 11.4

Hydrogen Bonding in H2O

• Ice has open lattice- like structure.

• Density Ice < liquid

• solid water floats on water.

Snow flake: www.snowcrystals.com

11.4

11.4 Intermolecular Forces

• Electrons in a molecule have some freedom to move about • molecule may have a nonuniform distribution of electron

density, resulting in temporary dipole— instantaneous dipole or induced dipole

An instantaneous dipole in one molecule can induce dipoles in neighboring molecules.

11.4 Intermolecular Forces

London dispersion forces or simply dispersion forces result attractive forces

As dispersion forces increase, intermolecular forces increase.

As the intermolecular forces increase, boiling points increase.

Polar Covalent Bonds: Electronegativity

Polarizability: Relative charge distribution, the electron cloud of an atom or molecule, is distorted from its normal shape by an external electric field, which may be caused by the presence of a nearby ion or dipole.

11.4

11.4 Intermolecular Forces

Ion-dipole interactions are Coulombic attractions between ions (either positive or negative) and polar molecules.

Magnitude of ion-dipole interactions depends on:

Ion’s the charge and the ion size Polar molecule’s dipole moment and size

Na+

Cl 

Summary of IMF in descending Order

Type of interaction Factors Responsible for Example

Interaction

Ion-Ion Charge of the ion, 400-4000 kJ/mol K+ and Cl-

Ion–dipole Ion charge, magnitude of dipole,

40-600 kJ/mol

K+ in H2O

Hydrogen bonding Very polar X-H bond to atom Y

(X,Y = F,N,O), 10-40 kJ/mol

H2O–H2O

Dipole–dipole Dipole moment, 5-25 kJ/mol CH2Cl2–CH2Cl2

Dipole-induced

dipole

Dipole moment of polar molecule &

polarizability of nonpolar molecule,

2-10 kJ/mol

O2–H2O

Induced dipole–

induced dipole (LDF)

Polarizability, 0.05-40 kJ/mol Br2–Br2

11.4

Interacting particles

Are polar

molecules

involved?

Are ions

involved?

Are polar

molecules

involved?

Are Hydrogen

atoms bonded to O,

N or F involved?

London

forces

(induced

dipoles)

Dipole–

dipole

forces

Hydrogen

bonding

Ion–dipole

forces

Ionic

bonding

No No

No Yes

Yes

Concept Map - IMFs

11.4

11.1

What kind(s) of intermolecular forces exist in (a) CCl4(l) (b) CH3COOH(l)

Draw Lewis dot structures and apply VSEPR theory to determine whether each molecule is polar or nonpolar.

(a) CCl4 is nonpolar, so the only intermolecular forces are dispersion forces.

(b) CH3COOH is polar and contains an O-H bond, so it exhibits dipole-dipole interactions (including hydrogen bonding) and dispersion forces.

• Particles are in constant motion

• Particles are in close contact

• Liquids are almost incompressible

• Liquids do not fill the container

• Intermolecular forces are relevant molecules are inter-dependent

2 1

Properties of Liquids11.1

• Energy needed to break the intermolecular bonds in liquids and solids

– Energy is released when intermolecular “bonds” are formed

• N is blue and H is white

What kind of change is represented? Deposition gs

Is energy absorbed or released when the change occurs? Released

2 2

Phase Changes11.1

11.1 Phase Changes

Phase Change (State Change): A change in physical form but NOT the chemical identity of a substance.

NO bonds are broken! Break IMF’s

Chapter 8/24

Enthalpies of Physical and Chemical Change

Enthalpy of Fusion (∆Hfusion opposite ∆Hfreezing): The amount of heat necessary to melt a substance without changing its temperature, s  l or l  s

∆Hfusion = - ∆Hfreezing

Enthalpy of Vaporization (∆Hvap opposite ∆Hcondensation): The amount of heat required to vaporize a substance without changing its temperature, l  g or g l

∆Hvaporization = - ∆Hcondensation

Enthalpy of Sublimation (∆Hsublimation opposite ∆Hdeposition): The amount of heat required to convert a substance from a solid to a gas without going through a liquid phase, s  g or g s

∆Hsublimation = - ∆Hdeposition

+40.7 kJ/mol -40.7 kJ/mol

+6.02 kJ/mol -6.02 kJ/mol

11.1

ΔHfusion (kJ/mol) Melting point (°C) ΔHvap (kJ/mol) Boiling point (°C)

Methane 0.94 –182.5 8.2 –161.6

Ethane 2.86 –182.8 14.7 –88.6

Propane 3.53 –187.6 19.0 –42.1

Methanol 3.16 –97.0 35.3 64.7

Ethanol 5.02 –114.3 38.6 78.4

1-Propanol 5.20 –127 41.4 97.2

Water 6.01 0.0 40.7 100.0

Na 2.60 97.82 97.42 881.4

NaBr 26.11 755 160.7 1390

Table 12.5.1 - Selected Heats of Fusion and Vaporization at the Temperature of the Normal Phase Transition

11.5

Phase Diagram summarizes the conditions (temperature and pressure) at which a substance exists as a solid, liquid, or gas.

Phase Diagrams Normal Boiling Point and Melting Point : Occurs at 1 atm.

Critical Point: A combination of temperature and pressure beyond which a gas cannot be liquefied.

• Critical Temperature: The temperature beyond which a gas cannot be liquefied regardless of the pressure.

• Critical Pressure: The pressure beyond which a liquid cannot be vaporized regardless of the temperature.

Supercritical Fluid: A state of matter beyond the critical point that is neither liquid nor gas.

Triple Point: A point at which three phases coexist in equilibrium.

11.1

Intermolecular Forces and the Liquid State

11.2 Vapor Pressure

• Vapor Pressure - Pressure of the liquid molecules in the vapor state

• For a liquid to vaporize, sufficient energy must be added (endothermic) to overcome the break IMFs

• Dynamic equilibrium: State in which rate of forward is equal to the rate of reverse

• Equilibrium vapor pressure: Pressure exerted by a vapor over a liquid in a closed container at a given temperature

Dynamic Equilibrium and Vapor Pressure

11.2

What is the chemical composition of the vapor?

• Liquid with the stronger IMFs has lower vapor pressure at a given T - Inverse relationship

Vapor Pressure and IMFs11.2

S=C=S

Table 11.2.1 - Vapor Pressure of Some Common Liquids

Boiling Point

N o

rm a

l b

p =

7 8

.5 °C

Water boils at 100°C at sea level

Sea level 1 atm = 760 mmHg

N o

rm a

l b

-40 -20 0 20 40 60 80 100 120 Temperature, °C

V a

p o

r p

re ss

u re

, m

m H

1000

500

b p

= 3

4 .6

°C

Salt Lake City (4400 ft) P=650 mmHg

Water boils at 95°C in Salt Lake City

11.2 • Temperature at which Pvap = Patm

– Normal boiling point: Patm = 760 mmHg (1 atm)

• As strength of IMF increases, the boiling point increases

• At reduced external pressure (higher altitude)

Boiling occurs at a lower temperature

Boiling Point and Baking11.2 Baking at Higher elevations (lower air pressure) is a SCIENCE!

Liquids evaporate faster, amounts of flour, sugar and liquids are changed to prevent batter that is too moist, dry or gummy

Gases expand more - doughs and baked goods rise faster easier for gas bubbles to rise. Leavening agents (baking soda and baking powder) are decreased

At elevations over 3500 feet, the oven temperature for batters and doughs should be 25oF higher than at sea level, WHY?

…leavening and evaporation are faster, use a higher temperature to set the structure before overexpansion and dry out

• IMFs affect- Enthalpy of vaporization, Vapor pressure, Boiling point of liquids

• For a series of liquids, as IMF strength increases:

– Energy needed to vaporize the liquids increases (ΔHvap increases)

– Liquid vapor pressure decreases

– Liquid boiling point increases

IMF Summary

Molecules in

the liquid

state

Hvap Volatility Equilibrium

vapor

pressure

Boiling

point

Strong IMFs More

endothermic

Low Low High

Weak IMFs Less

endothermic

High High Low

11.2

• Mathematical relationship between vapor pressure (P), temperature (T), and strength of IMFs (related to ΔHvap)

R - Ideal gas constant (R = 8.314×10–3 kJ/K·mol)

Clausius–Clapeyron Equation

vap

- = +

= + x

y m

In 1

TR P

H C

     

11.2

Clausius-Clapeyron Equation Measured values to obtain the relationship between VP & T:

NOTE: T must be in Kelvin! R = 8.3145 J/mol-K

Plot ln P versus 1/T

slope = – ΔHvap R

Ethanol

11.2

Clausius-Clapeyron Equation What is the ΔHvap for diethyl ether, CH3CH2OCH2CH3 if P1 = 57.0 mm Hg at T1= -22.8

oC and P2 = 534 mm Hg at T2= 25.0oC ?

ln = –P2 P1

1 T2

1 T1

-ΔHvap R

Ln = – 534

57.0

1 250.35

1 298.15

-ΔHvap 8.315 J K-1mol-1

ΔHvap = 29048.8 J/mol = 29.0 kJ/mol

11.2

Intermolecular Forces and the Liquid State

11.3 Other Properties of Liquids

11.3 Properties of Liquids

Surface Tension, measure of the elastic force, energy required to stretch or increase the surface of a liquid by a unit area (1 cm2).

Measure of force required to "break" the surface of a liquid, “skin”

A liquid with strong intermolecular forces has a high surface tension.

Molecules on the surface are not pulled upward but are pulled downward and sideways, causes tightening of surface molecules

http://www.npr.org/blogs/krulwich/2013/04/21/17794960 5/a-wet-towel-in-space-is-not-like-a-wet-towel-on-earth

http://www.npr.org/blogs/krulwich/2013/04/21/177949605/a-wet-towel-in-space-is-not-like-a-wet-towel-on-earth

11.3 Properties of Liquids

Viscosity, is a measure of a fluid’s resistance to flow, units of N · s/m2,

Higher the viscosity, the more slowly a liquid flows.

Liquids that have strong IMFs have higher viscosities than those that have weaker IMFs.

Higher Temperature = Lower viscosity

• Ek > Ebarrier required to move past another molecule.

Longer molecule = Higher viscosity • Flexible, longer molecules become tangled and hinder flow

Capillary action - liquid is pulled up into a narrow glass cylinder

A competition between:

• Adhesion: attractions between unlike molecules.

• Cohesion: attractions between like molecules.

Shape of the meniscus illustrates the relative strength of the adhesive and cohesive forces

11.3 Properties of Liquids

Adhesion < Cohesion Convex meniscus is formed

Adhesion > Cohesion Concave meniscus is formed

H2O adheres to the wall

Intermolecular Forces and the Liquid State

11.5 Intermolecular Forces and the Properties of Liquids

Table 11.5.1 - Properties of Some Common Nonpolar Species

Compound

Molar mass

(g/mol) ΔHvap (kJ/mol) Boiling point (°C)

He 4.0 0.08 –268.9

Ne 20.3 1.7 –246.1

N2 28.0 5.6 –195.8

O2 32.0 6.8 –183.0

Ar 39.9 6.4 –185.9

Cl2 70.9 20.4 –34.0

Br2 159.8 30.0 58.8

11.5

• Weaker the IMF the lower the vapor pressure

• Consider a series of atoms and molecules that have only dispersion forces

– As the molar mass increases, enthalpy of vaporization increases

– Larger the molecule, greater the number of electrons, and greater the polarizability

– Greater the surface area available for contact, greater the dispersion forces

Enthalpy of Vaporization11.5

• If the molar masses are similar:

– Molecule with the greater IMF will have the higher boiling point

• In each pair below, the molecule that is polar has the higher boiling point

Quantitative Comparison of IMFs11.5

Nonpolar Polar Molar mass

(g/mol)

Boiling point

(°C)

Molar mass

(g/mol)

Boiling point

(°C)

N2 28 –196 CO 28 –192

SiH4 32 –112 PH3 34 –88

GeH4 77 –90 AsH3 78 –62

Br2 160 59 ICl 162 97

Summary of IMF in descending Order

Type of interaction Factors Responsible for Example

Interaction

Ion-Ion Charge of the ion, 400-4000 kJ/mol K+ and Cl-

Ion–dipole Ion charge, magnitude of dipole,

40-600 kJ/mol

K+ in H2O

Hydrogen bonding Very polar X-H bond to atom Y

(X,Y = F,N,O), 10-40 kJ/mol

H2O–H2O

Dipole–dipole Dipole moment, 5-25 kJ/mol CH2Cl2–CH2Cl2

Dipole-induced

dipole

Dipole moment of polar molecule &

polarizability of nonpolar molecule,

2-10 kJ/mol

O2–H2O

Induced dipole–

induced dipole (LDF)

Polarizability, 0.05-40 kJ/mol Br2–Br2

11.5

Interacting particles

Are polar

molecules

involved?

Are ions

involved?

Are polar

molecules

involved?

Are Hydrogen

atoms bonded to O,

N or F involved?

London

forces

(induced

dipoles)

Dipole–

dipole

forces

Hydrogen

bonding

Ion–dipole

forces

Ionic

bonding

No No

No Yes

Yes

Concept Map - IMFs

11.5