PowerPoint to accompany Chapter 10 Intermolecular Forces, Liquids and Solids

PowerPoint to accompany

Chapter 10

Intermolecular Forces, Liquids

and Solids

Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia

States of Matter,Forces & Equations of State

The fundamental difference between states of matter is the distance between particles.

Figure 10.1

(Kinetics, Heat, Gas Laws) versus (Polar, Hydrogen Bonding, Ionic, Coulomb’s Law, Crystal Field Energy, Berthelot’s for Liquids & Solids)


The States of Matter: Gas, Liquid & Solid

Table 10.1

The state a substance is in at a particular temperature and pressure depends on two antagonistic entities: The kinetic energy of the

particles The strength of the

attractions between the particles

Disorder-order kas a kinetic effect , it takes time to untangle a knot!


Intermolecular Forces: Coulomb’s Law & Instantaneous Dipoles

The attractions between molecules are not nearly as strong as the intramolecular attractions that hold compounds together.

Figure 10.2


Intermolecular Forces

These intermolecular forces are referred to as van der Waals forces or Van der Waals Bonds.

Figure 10.2


They are, however, strong enough to control physical properties such as boiling and melting points, vapour pressures and viscosities, freezing point depression.

Figure 10.2

Intermolecular Forces: Dipole-Dipole, Hydrogen Bonding [H-(N,O,F)], London Dispersion, Crystal Lattice Energy


Van der Waals Forces: Gas Liq. SolidDipole-dipole forces: condensation of water,

ammonia, freons (Rain, refrigeration cycles) Hydrogen bonding: R-(OH)--(OH), (NH),

(FH), cohesion, adhesion, glues, mud, food

Ion-dipole forces: water and ions, solvation energy: K+ or Cl- w 6-7 H2O’s (ocean, blood)

London dispersion forces: induced polarizations (liquid nitrogen, gas or unlike substance dissolution)


Dipole-Dipole Interactions: condensation, viscosity, AC conductivity

Molecules that have permanent dipoles are attracted to each other:

The positive end of one is attracted to the negative end of the other and vice-versa.

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

Heat and disorder hold these forces at bay

Figure 10.4


Dipole-Dipole Interactions: Boiling Point and Freezing (Melting) Point

The more polar the molecule, the higher (warmer) is its boiling point. (It takes more heat)

The Larger, Heavier & more complicated molecules also have higher boiling points and warmer freezing/melting points too.

Table 10.2


Ion-Dipole Interactions

Ion-dipole interactions are an important force in solutions of ions. Na+ and Cl- typically have 6-7 waters in their solvation shell.

The strength of these forces is what makes it possible for ionic substances to dissolve in polar solvents.

Figure 10.3


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), they occasionally wind up on the same side of the atom.



At that instant, the helium atom is polar, with an excess of electrons on the left side and a shortage on the right side. Call this an instantaneous dipole if you like but this will occur by induction from a neighbouring He atom whether it is lopsided or not, due to insufficient screening of 2 protons by opposite electrons.


London Dispersion Forces: induced dipoles

Another helium nearby would then have a dipole induced in it, as the electrons on the left side of helium atom 2 repel the electrons in the cloud on helium atom 1. Electrical induction occurs at high density, low energy.

Figure 10.5


London Dispersion Forces: induced dipoles

Figure 10.5

London dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and an induced dipole. Helium is small, symmetric, electron velocities are fast. Induced dipoles oscillate so Liquid He displays superfluid properties. http://www.youtube.com/watch?v=2Z6UJbwxBZI



These forces are present in all molecules and atoms, whether they are polar or nonpolar.

The tendency of an electron cloud to distort in this way is called polarizability.

The more polar, the larger, the more uneven bond types and larger atoms with d, f orbitals are all selectively more polarizable.

Figure 10.5


Molecular Shape The shape of the molecule

affects the strength of dispersion forces: long, skinny molecules (like n-pentane) tend to have stronger dispersion forces than short, fat ones (like neopentane).

This is due to the increased surface area in n-pentane which makes it have greater LDS and easier to stack.

Figure 10.6


Factors Affecting London Forces

The strength of dispersion forces tends to increase with increased molecular weight.

Larger atoms have larger electron clouds which are easier to polarize.

Table 10.3


Which Have a Greater Effect:Dipole-Dipole Interactions or Dispersion Forces? If molecules 2 different compounds are of

comparable size and shape then their LDS forces will be comparable. Thus, differences in physical properies likely depend on dipole-dipole interactions. e.g. Methyl Iodide CH3I has dipole moment 1.62 D and TB = 316K while Acetomitrile has dipole moment of 3.9 D and TB = 354.8K

From these data we see that the heavier Methyl iodide at 142g/Mol depends mainly on LDS while Acetonitrile at 142 g/Mol mainly relies on dipoles.


Which Have a Greater Effect:Dipole-Dipole Interactions or Dispersion Forces? In any compound, both forces may contribute to

behaviour. e.g. HCl depends 70% on LDS for its behaviour while only 30% is due to its dipole moment of 1.05 D. When dipole moments are moderate, both dipoles and LDS play a role.

If one molecule/atom is much larger than another, dispersion forces will likely determine its physical properties, e.g. TB Xe = 166K while He is 4.6 K.


Which Have a Greater Effect:Dipole-Dipole Interactions or Dispersion Forces? Practice: Among Br2 , Ne, HCl, HBr and N2 … A) Which has the greatest Dipole Dipole forces?

B) Which has the greatest London Dispersion Forces?


Which Have a Greater Effect:Dipole-Dipole Interactions or Dispersion Forces? Practice: Among Br2 , Ne, HCl, HBr and N2 … A) Which has the greatest Dipole Dipole forces?

HCl is the smallest dipole with greatest dipole forces because H is the same and Cl is smaller and more electronegative 3.0 versus 2.8 for Br. The others aren’t polar. Dipoles go as Qr.

B) Which has the greatest London Dispersion Forces? Br2 is the largest mass 159 g/Mol thus most polarizable


How Do We Explain These Trends in Boiling points @ 1bar?

Chalcogen (Group 16) hydrides: The polar series follows the trend from H2Te through H2S, but H2O is anomalous. (Strong H-bonds w/O)

(Group 14) hydrides: The nonpolar series (SnH4 to CH4) follows the expected trend, but not Methane. (Small LDS)

Figure 10.7


Hydrogen Bonding: H w/ N, O or F

The dipole-dipole interactions experienced when H is bonded to tiny Period 2 elements at the right hand end of the P block: N, O, or F are unusually strong.

We call these interactions hydrogen bonds.

Figure 10.8


Hydrogen Bonding

Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine.

Also, when hydrogen is bonded to one of those very electronegative

elements, the naked proton = hydrogen nucleus is exposed.


Summarising Intermolecular Forces

Figure 10.12


Intermolecular Forces Affect Many Physical Properties

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

Figure 10.15


Viscosity: shear stress/shear strain (resistance to flow or deformation)

Resistance of a liquid to flow is called viscosity.

It is related to the ease with which molecules can move past each other.

Viscosity increases with stronger intermolecular forces and decreases with higher temperature. Fig. 10.13

Table 10.4


Surface Tension: Liquid Cohesion

Surface tension results from the net inward force experienced by the molecules on the surface of a liquid.

Figure 10.14

Figure 10.15


Phase Changes: Boiling etc.

Figure 10.17


Energy Changes Associated with Changes of State

Heat of Fusion: Energy required to change a solid at its melting point to a liquid.

Heat of Vaporization: Energy required to change a liquid at its boiling point to a gas.

Figure 10.18


Energy Changes Associated with Changes of State

The heat added to the system at the melting and boiling points goes into pulling the molecules further apart from each other.

The temperature of the substance does not rise during the phase change.

Figure 10.19


Vapour Pressure At any temperature, some molecules in a liquid have

enough energy to escape.

As the temperature rises, the fraction of molecules that have enough energy to escape increases.

Figure 10.21


Vapour Pressure

As more molecules escape the liquid, the pressure they exert increases.

The liquid and vapour reach a state of dynamic equilibrium: liquid molecules evaporate and vapour molecules condense at the same rate.

Figure 10.20


Boiling Point

The boiling point of a liquid is the temperature at which its vapour pressure equals atmospheric pressure.

The normal boiling point is the temperature at which its vapour pressure is 101.3 kPa.

Figure 10.22


Phase DiagramsPhase diagrams display the state of a substance at various pressures and temperatures and the places where equilibria exist between phases.

Figure 10.24


Phase Diagrams The AB line is the liquid-vapour interface. It starts at the triple point (A), the point at which

all three states are in equilibrium.

Figure 10.24


Phase DiagramsIt ends at the critical point (B); above this critical temperature and critical pressure, the liquid and vapour are indistinguishable from each other.

Figure 10.24


Phase Diagrams

Each point along this line is the boiling point of the substance at that pressure.

Figure 10.24


Phase Diagrams The AD line is the interface between liquid and

solid. The melting point at each pressure can be found

along this line.

Figure 10.24


Phase Diagrams Below A, the substance cannot exist in the liquid

state. Along the AC line, the solid and gas phases are

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

Figure 10.24


Phase Diagram of Water

Note the high critical temperature and critical pressure: These are due to the

strong van der Waals forces between water molecules.

Figure 10.25


Phase Diagram of Water

The slope of the solid–liquid line is negative. This means that

the melting point decreases with increasing pressure.

Figure 10.25


Phase Diagram of Carbon Dioxide

Carbon dioxide cannot exist in the liquid state at pressures below 518 kPa; CO2 sublimes at normal pressures.

Figure 10.25


Phase Diagram of Carbon DioxideThe low critical temperature and critical pressure for CO2 make supercritical CO2 a good solvent for extracting nonpolar substances (such as caffeine).

Figure 10.25


Structures of Solids

We can think of solids as falling into two groups:

i) crystalline

particles are in a highly ordered arrangement.

ii) amorphous

no particular order in the arrangement of particles.

Figure 10.28


Crystalline Solids

Because of the order in a crystal, we can focus on the repeating pattern of arrangement called the unit cell.

Figure 10.30

Figure 10.31


Crystalline Solids

There are several types of basic arrangements in crystals, such as the ones shown above.

Figure 10.32


Close Packing of Spheres The structures adopted by crystalline solids are

those that bring particles into closest contact to maximise the attractive forces between them.

Consider how equal-sized spheres can pack with the minimum amount of empty space.

Figure 10.35


Bonding in Solids

Table 10.7


Covalent-Network Solids

Diamonds are an example of a covalent-network solid in which atoms are covalently bonded to each other. They tend to be hard and have high melting points.


Ionic Solids

(a) Green: chlorine; Grey: cesium

(b) Yellow: sulfur; Grey: zinc

(c) Green: calcium; Grey: fluorine

Figure 10.40


Metallic Solids

Metals are not covalently bonded, but the attractions between atoms are too strong to be van der Waals forces.

In metals, valence electrons are delocalised throughout the solid. Figure 10.41

Documents

PowerPoint to accompany Chapter 10 Intermolecular Forces, Liquids and Solids