chapters 11&12 intermolecular forces liquids and solids

CHEM 1311 Dr. Pahlavan

1

Chapters 11&12

Intermolecular Forces

Liquids and Solids The physical properties of a substance depends upon its physical state. Water vapor, liquid water and ice all

have the same chemical properties, but their physical properties are considerably different.

a) A collection of widely separated molecules.

b) The kinetic energy of the molecules is greater than any attractive forces between the molecules.

c) The lack of any significant attractive force between molecules allows a gas to expand to fill its

container.

d) If attractive forces become large enough, then the gases exhibit non-ideal behavior.

Liquids properties

a) The intermolecular attractive forces are strong enough to hold molecules close together.

b) Liquids are more dense and less compressible than gasses.

c) Liquids have a definite volume, independent of the size and shape of their container.

d) The attractive forces are not strong enough, however, to keep neighboring molecules in a fixed position

and molecules are free to move past or slide over one another.

Thus, liquids can be poured and assume the shape of their containers

Solids properties

a) The intermolecular forces between neighboring molecules are strong enough to keep them locked in

position.

b) Solids (like liquids) are not very compressible due to the lack of space between molecules.

c) If the molecules in a solid adopt a highly ordered packing arrangement, the structures are said to be

crystalline. d) Due to the strong intermolecular forces between neighboring molecules, solids are rigid.

e) The state of a substance depends on the balance between the kinetic energy of the individual particles.

f) (molecules or atoms) and the intermolecular forces . Intermolecular forces try to draw the particles

together.

g) Kinetic energy keeps the molecules apart and moving around, and is a function of the temperature of

the substance.

Gases have weaker intermolecular forces than liquids. Liquids have weaker intermolecular forces than

solids. Solids and liquids have particles that are fairly close to one another, and are thus called "condensed

phases" to distinguish them from gases.

Changing the state of a substance

Temperature - Heating and cooling can change the kinetic energy of the particles in a substance, and so, we

can change the physical state of a substance by heating or cooling it. Cooling a gas may change the state to a

liquid and cooling a liquid may change the state to a solid.


2

Pressure - Increasing the pressure on a substance forces the molecules closer together, which increases the

strength of intermolecular forces. Increasing the pressure on a gas may change the state to a liquid. Increasing

the pressure on a liquid may change the state to a solid.

Intermolecular Forces- Intermolecular forces are generally much weaker than covalent bonds. Only 16 kJ/mol

of energy is required to overcome the intermolecular attraction between HCl molecules in the liquid state (i.e.

the energy required to vaporize the sample). However, 431 kJ/mol of energy is required to break the covalent

bond between the H and Cl atoms in the HCl molecule.

Thus, when a molecular substance changes states the atoms within the molecule are unchanged.

The temperature at which a liquid boils reflects the kinetic energy needed to overcome the attractive

intermolecular forces (likewise, the temperature at which a solid melts).

Thus, the strength of the intermolecular forces determines the physical properties of the substance.

Attractive forces between neutral molecules are;

a) Dipole-dipole forces b) London dispersion forces c) Hydrogen bonding forces

Typically, dipole-dipole and dispersion forces are grouped together and termed van der Waals forces

(sometimes the hydrogen bonding forces are also included with this group).

Attractive forces between neutral and charged (ionic) molecules are;

a) ion-dipole forces - Note that all of these forces will be electrostatic in nature

b) ion-dipole - Involves an interaction between a charged ion and a polar molecule (i.e. a molecule with a

dipole).

Cations are attracted to the negative end of a dipole and anions are attracted to the positive end of a dipole.

The magnitude of the interaction energy depends upon the charge of the ion (Q), the dipole moment of the

molecule (u) and the distance (d) from the center of the ion to the midpoint of the dipole.

E α ( Q u / d2 )

Ion-dipole forces are important in solutions of ionic substances in polar solvents (e.g. a salt in aqueous solvent)


3

Dipole-Dipole Forces - A dipole-dipole force exists between neutral polar molecules. Polar molecules attract

one another when the partial positive charge on one molecule is near the partial negative charge on the other

molecule. The polar molecules must be in close proximity for the dipole-dipole forces to be significant

Dipole-dipole forces are characteristically weaker than ion-dipole forces. Dipole-dipole forces increase with an

increase in the polarity of the molecule.

Boiling points increase for polar molecules of similar mass, but increasing dipole:

London Dispersion Forces - Nonpolar molecules

would not seem to have any basis for attractive

interactions.

However, gases of nonpolar molecules can be liquefied

indicating that if the kinetic energy is reduced, some

type of attractive force can predominate.

Hydrogen Bonding - a strong intermolecular force—

the hydrogen bond—forms when a H atom covalently

bonded to an O, a N, or a F atom of one molecule is

simultaneously attracted to an O, a N, or a F atom of

another molecule. Hydrogen bonding accounts for some of the unusual properties of water (for example, an

unusually high boiling point and the fact that the density of liquid water is greater than the density of ice).

Hydrogen bonding also governs aspects of the behavior of biological molecules such as proteins and the nucleic

acids.

They are very important in the organization of biological molecules, especially in influencing the structure of

proteins.

δ- δ- …….. H O …… H O ………H

δ+ δ+

H δ+ H δ+

Water is unusual in its ability to form an extensive hydrogen bonding network. As a liquid the kinetic energy of

the molecules prevents an extensive ordered network of hydrogen bonds. When cooled to a solid the water

molecules organize into an arrangement which maximizes the attractive interactions of the hydrogen bonds.

This arrangement of molecules has greater volume (is less dense) than liquid water, thus water expands when

frozen.

Properties of Liquids I ) Viscosity II) Surface Tension

Viscosity - The resistance of a liquid to flow is called its viscosity. The greater the viscosity, the more slowly it flows. Measuring viscosity – Determines the time that a liquid takes to flow out of a pipette under the force of gravity

and how fast an object (steel ball) sinks through the liquid under gravitational force.

Substance Molecular

Mass (amu)

Dipole

moment, u

(D)

Boiling

Point

(°K)

Propane 44 0.1 231

Dimethyl

ether 46 1.3 248

Methyl

chloride 50 2.0 249

Acetaldehyde 44 2.7 294

Acetonitrile 41 3.9 355


4

The Physical Basis of Viscosity - Viscosity is a measure of the ease with which molecules move past one

another. It depends on the attractive force between the molecules and also it depends on whether there are

structural features which may cause neighboring molecules to become "entangled".

Viscosity decreases with increasing temperature - the increasing kinetic energy overcomes the attractive forces

and molecules can more easily move past each other.

Surface Tension - By definition the molecules of a liquid exhibit intermolecular attraction for one

another.

What happens to molecules at the surface in comparison to those in the interior of a liquid?

Molecules in the interior experience an attractive force from neighboring molecules which surround on all sides.

Molecules on the surface have neighboring molecules only on one side (the side facing the interior) and thus

experience an attractive force which tends to pull them into the interior. The overall result of this asymmetric

force on surface molecules is that:

a) The surface of the liquid will rearrange until the least number of molecules are present on the surface.

b) In other words the surface area will be minimized and a sphere has the smallest surface area to volume

ratio.

c) The surface molecules will pack somewhat closer together than the rest of the molecules in the liquid

d) The surface molecules will be somewhat more ordered and resistant to molecular disruptions.

Thus, the surface will seem to have a "skin". The "inward" molecular attraction forces, which must be overcome

to increase the surface area, are termed the "surface tension".

Surface tension is the energy required to increase the surface area of a liquid by a unit amount.

Water : Intermolecular hydrogen bonds and surface tension at 20°C is 7.29 x 10-2

J/m2

Mercury : Intermolecular metallic (electrostatic) bonds and surface tension at 20°C is 4.6 x 10-1

J/m2

Cohesive forces bind molecules of the same type together and Adhesive forces bind a substance to a surface

For example, attractive forces (hydrogen bonding) exists between glass materials (Silicon dioxide) and water.

This is the basis of "capillary" action, where water can move up a thin capillary, against the force of gravity.

Surface tension "pulls" neighboring water molecules along. The liquid climbs until the adhesive and cohesive

forces are balance by the force of gravity.

Changes of State - The three states of matter include , solid , liquid , and gas . In general, matter in one state can be changed into either of the other two states. Such transformations are called "phase changes". Each change of state is accompanied by a change in the energy of the system.

Whenever the change involves the disruption of intermolecular forces, energy must be supplied . The disruption

of intermolecular forces accompanies the state going towards a less ordered state. As the strengths of the

intermolecular forces increase, greater amounts of energy are required to overcome them during a change in

state. The melting process for a solid is also referred to as fusion.


5

The enthalpy change associated with melting a solid is often called the heat of fusion (Hfus) ,

Ice Hfus = 6.01 kJ/mol = Hf

The heat needed for the vaporization of a liquid is called the heat of vaporization (Hvap),

Water Hvap = 40.67 kJ/mol = Hv

Less energy is needed to allow molecules to move past each other than to separate them totally.

Vaporization requires the input of heat energy. Refrigerators use the evaporation of Freon (CCl2F2) to remove

heat inside the fridge. The Freon is condensed outside the cabinet (usually in coils at the back) in a process

which releases heat energy (the coils will be warm).

Heating Curves – The heating of ice at -25 °C to +125 °C at constant pressure (1 atm) will exhibit the following characteristics. Initially, the heat input is used to increase the temperature of the ice, but the ice does not change phase (remains a solid). As the temperature approaches some critical point (i.e. the melting temperature of ice), the kinetic energy of the

molecules of water is sufficient to allow the molecules to begin sliding past one another.

As the ice begins to melt, additional input of heat energy does not raise the temperature of the water, rather it is

used to overcome the intermolecular attraction during the phase change from solid to liquid.

Once the water is in a liquid phase, increasing the amount of heat input raises the temperature of the liquid

water.

As the temperature approaches another critical point (the vaporization, or boiling, temperature of water) the

kinetic energy of the molecules is sufficient to allow the separation of molecules into the gas phase.

As the liquid begins to boil. Additional input of heat energy does not raise the temperature of the water, rather it

is used to overcome the intermolecular attractions during the phase change from liquid to gas. Once the water is

in the gas phase, additional heat input raises the temperature of the water vapor.

Note: greater energy is needed to vaporize water than to melt it.

Heating ice, water and water vapor - In the region of the curve where we are not undergoing a phase

transition, we are simply changing the temperature of one particular phase of water (either solid, liquid or gas)

as a function of heat input. The slope of the lines relates temperature to heat input. The greater the slope, the

greater the temperature change for a given unit of heat input.

The amount of heat needed to change the temperature of a substance is given by the specific heat or molar heat

capacity.

Specific heat of ice = 2.09 J/g K = 0.50 cal/g K , Specific heat of water = 4.18 J/g K = 1.00 cal/g K

Specific heat of water vapor = 1.84 J/g K


6

In the regions of the curve where we are undergoing a phase transition, the heat energy input is not raising the

temperature of the sample, rather it is being used to disrupt the intermolecular forces.

Hfus = 6.01 kJ/mol and Hvap = 44.0 kJ/mol

Example 1 - Calculate the enthalpy change for converting 2 moles of ice at -25°C to +125°C.

Converting to grams: (2 mol).(18 g/mol) = 36 g

Heating ice from -25 to 0°C: (25K).(2.09 J/g K).(36 g) = 1.88 kJ

Fusion of ice to liquid water: (6.01 kJ/mol).(2 mol) = 12.0 kJ

Heating of water from 0 to 100°C: (100K).(4.18 J/g K).(36 g) = 15.1 kJ

Vaporization of water to water vapor: (44.0 kJ/mol).(2 mol) = 88.0 kJ

Heating of water vapor from 100 to 125°C: (1.84 J/g K).(25K).(36 g) = 1.66 kJ

Grand total: 1.88 + 12.0 + 15.1 + 88.0 + 1.66 = 119 kJ

Alternate Solution

H = Mice . Cice . T + Mice .Hf + Mw . Cw . T + Mw . Hv + Mv . Cv . T

= (2x18)(2.09)( 0 +25) + (2)(6.01) + (2x18)(4.18)(100 -0) + (2)( 44.0) + (2x18)( 1.84)(125-100)

= 1.88 kJ + 12.0 kJ + 15.1 kJ + 88.0 kJ + 1.66 kJ = 119 kJ

Critical Temperature (Tc) and Pressure (Pc) - Gases can be liquified by either decreasing the temperature or

increasing the pressure. As long as the temperature is not too high, we can use pressure to liquefy a gas. As

temperatures increase it becomes more difficult to use pressure to liquefy a gas (due to the increasing kinetic

energy).

For every substance there is a temperature above which it is impossible to liquefy the gas regardless of the

increase in pressure.

The highest temperature at which a substance can exist as a liquid is called its critical temperature.

The critical pressure is the pressure required to bring about condensation at the critical temperature.

For example, oxygen has a critical temperature of 154.4 K. It cannot be liquefied until the temperature is

reduced to this point. At this temperature, the pressure needed to liquefy oxygen is 49.7 atm.

Vapor Pressure - Suppose we have a closed container into which we pour some water. As soon as we add the

water we check a pressure gauge connected to the container. We let the container sit for a while and then we

check the pressure again. What might the pressure guage indicate?

As the water evaporates the pressure exerted by the vapor above the liquid increases, until at some point, the

pressure reaches a constant value called the vapor pressure of the substance.

(The molecular basis of vapor pressure)


7

The kinetic energy of the molecules at the surface of a liquid varies over a range of values.

Some of the molecules have enough kinetic energy to overcome the attractive forces between the molecules.

The weaker the attractive forces, the greater the fraction of molecules with enough kinetic energy to escape.

The greater the fraction of molecules which can escape the liquid, the greater the vapor pressure .

Not only can water molecules leave the surface, but molecules in the vapor phase can also hit and go into the

water.

Initially, there are no molecules in the vapor phase and the number of molecules in the vapor which are

rejoining the water is zero. As time goes on there are more molecules in the vapor phase and the number of a

vapor molecule striking the water increases.

At some point in time the number of vapor molecules rejoining the water equals the number leaving to go into

the vapor phase. an equilibrium has been reached, and the pressure has stabilized at the characteristic vapor

pressure of the substance.

Vapor pressure increases with temperature. At higher temperature more molecules have the necessary

kinetic energy to escape the attractive forces of the liquid phase. The more molecules in the vapor phase, the

higher the vapor pressure.

What if molecules in the interior of the liquid decides to leave the liquid phase and go into the vapor phase?

This interior bubble will rapidly collapse if the external pressure is greater than the vapor pressure.

If the external pressure is equal to, or lower than the vapor pressure, then the bubble will remain or expand and

the liquid boils.

Vapor pressure increases with increasing temperature. At 100°C the vapor pressure of water is 760 torr (1 atm)

or equal to the atmospheric pressure on the liquid (in an open container). At this temperature, interior bubbles

will not collapse and the water boils. At high altitudes (i.e. up in the Mountains) the air pressure is less than at

sea level. Thus, water will boil at a lower temperature (the vapor pressure needed to support a bubble is lower at

high altitude). Therefore, cooking times are longer for things that need to be boiled (e.g. boiled eggs take longer

to cook at high altitudes).


8

Phase Diagrams - Equilibrium can exist not only between the liquid and vapor phase of a substance but also

between the solid and liquid phases, and the solid and gas phases of a substance. A phase diagram is a graphical

way to depict the effects of pressure and temperature on the phase of a substance:

Phase Changes Involving Solids—the conversion of a solid to a liquid is called either fusion or melting; the

temperature at which this change occurs is the melting point. The quantity of heat required to melt a given

amount of a solid is the enthalpy (heat) of fusion. A plot of temperature versus time as a solid is slowly heated

is known as a heating curve; a similar plot for a liquid that is slowly cooled is known as a cooling curve. In

some cases, it is possible to cool a liquid below its freezing point without having a solid form, a process known

as supercooling. The conversion of a solid directly to a gas (vapor) is called sublimation. A plot of the vapor

pressure of a solid versus temperature is known as a sublimation curve. The quantity of heat required to convert

a given amount of solid directly to a gas is the enthalpy (heat) of sublimation, which is equal to the sum of the

enthalpy of fusion and the enthalpy of vaporization..

The "triple point" is the particular condition of temperature and pressure where all three physical states are in

equilibrium. Regions not on a line represent conditions of temperature and pressure where only one particular

phase is present. Gases are most likely under conditions of high temperature and solids are most likely under

conditions of high pressure.

Phase Diagram for Water - The frozen state of water (ice) is actually less dense than the liquid state, thus, the

liquid state is more compact than the solid state. Increasing pressure, which will favor compactness of the

molecules, will thus favor the liquid state.Increasing pressure will thus lower the temperature at which the solid

will melt. The melting curve slopes to the left, unlike most compounds.


9

At 100 °C the vapor pressure of water is 760 torr or 1 atm, thus at this temperature water will boil if it is at 1

atm of pressure. At pressures below 4.58 torr, water will be present as either a gas or solid, there can be no

liquid phase.

Structures of Solids

a) Crystalline solids - The atoms, molecules or ions pack together in an ordered arrangement. Such solids

typically have flat surfaces, with unique angles between faces and unique 3-dimensional shape.

Examples of crystalline solids include diamonds, and quartz crystals(graphite).

b) Amorphous solids - No ordered structure to the particles of the solid. No well defined faces, angles or

shapes. Often are mixtures of molecules which do not stack together well, or large flexible molecules.

Examples would include glass and rubber .

The Structure of Crystals—the structure of a crystal, which is a solid substance composed of an extended

periodic array of atoms, is described by a three-dimensional pattern called a lattice. There are three types of

cubic lattices: simple cubic cell, body-centered cubic (bcc), and face-centered cubic (fcc).

A unit cell, when repeated in three dimensions, generates the entire crystal lattice. Unit cell properties and

dimensions, often measured by X-ray crystallography, can be used to find atomic radii and the densities of

crystalline substances. The crystal structures of metals can be described as the packing of spheres. The two

types of close-packing of spheres, hexagonal close-packed (hcp) and cubic close-packed (ccp), both reduce

to the smallest possible fraction the volume occupied by holes or voids. A model in which cations occupy the

voids present among a close-packed array of anions generally works well for the crystal structures of ionic

substances.

Since the crystal is made up of an arrangement of identical unit cells, then an identical point on each unit cell

represents an identical environment within the crystal . The array of these identical points is termed the crystal

lattic .

The unit cells shown are cubic, all sides are equal length , and all angles are 90° . The unit cell need not be

cubic. The unit cell lengths along the x,y, and z coordinate axes are termed the a, b and c unit cell dimensions

The unit cell angles are defined as: a) the angle (α ) formed by the b and c cell edges b) the angle (β) formed

by the a and c cell edges c) the angle (γ) formed by the a and b cell edges.


10

The crystal structure of sodium chloride;

The unit cell of sodium chloride is cubic, and this is reflected in the shape of NaCl crystals. The unit cell can be

drawn with either the Na+ ions at the corners, or with the Cl

- ions at the corners.

If the unit cell is drawn with the Na+ ions at the corners, then Na

+ ions are are also present in the center of each

face of the unit cell. If the unit cell is drawn with the Cl- ions at the corners, then Cl

- ions are are also present in

the center of each face of the unit cell. Within the unit cell there must be an equal number of Na+ and Cl

- ions.

Example - for the unit cell with the Cl- ions at the center of the faces.

The top layer has (1/8+1/8+1/8+1/8+1/2)=1 Cl- ion, and (1/4+1/4+1/4+1/4)=1 Na

+ ion

The middle layer has (1/2+1/2+1/2+1/2)=2 Cl- ions and (1/4+1/4+1/4+1/4+1)=2 Na

+ ions

The bottom layer will contain the same as the top or 1 each Cl- and Na

+ ions. The unit cell has a total of 4 Cl

-

and 4 Na+ ions in it. This equals the empirical formula NaCl.

Coordination Number - The coordination number is the number of particles surrounding a particle in the

crystal structure.

In each packing arrangement above (hexagonal close pack, cubic close pack), a particle in the crystal has a

coordination number of 12. The NaCl (face centered cubic) has a coordination number of 6.

Bonding in Solids - Molecular Solids - Consist of atoms or molecules held together by intermolecular forces

(dipole-dipole, dispersion and hydrogen bonds) . These forces are weaker than chemical (covalent) bonds.

Therefore molecular solids are soft, and have a generally low melting temperature. Most substances that are

gasses or liquids at room temperature form molecular solids at low temperature (e.g. H2O, CO2).

Covalent Network Solids - in some solids, covalent bonds extend throughout a crystal. Such solids, known as

network covalent solids, generally are harder, have much higher melting points, and are less volatile than other

molecular solids. Diamond and graphite, two allotropes of carbon, are examples of these solids.

Ionic Solids - Held together by ionic bonds. The strength of the ionic interactions depends on the magnitude of

the charge of the ions. Thus, NaCl (single charge on both ions) has a melting point of 801°C, whereas MgO (2+,

2- charge on the ions) has a melting point of 2852°C.

Metallic Solids - Consist entirely of metal atoms. Typically hexagonal close packed, cubic close packed or

body-centered cubic structures. These have coordination numbers of either 12 or 8.Bonding is due to valence

electrons which are delocalized throughout the entire solid.Bonding is stronger than simple dispersion forces,

but there are insufficient electrons to form ordinary covalent bonds. The strength of the bonding increases with

the number of electrons available for bonding.

Delocalization of electrons is the physical basis for the ability of metals to carry electrical current (electrons are

free to move about the metal structure). The nucleus and inner core of electrons are in a "sea" of delocalized,

mobile valence electrons.

chapters 11&12 intermolecular forces liquids and solids

Documents