© 2009, prentice-hall, inc. chapter 11 intermolecular forces, liquids, and solids

© 2009, Prentice-Hall, Inc.

Chapter 11Intermolecular Forces,

Liquids, and Solids


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.


The States of Matter

• 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.


Intermolecular Forces

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



They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities.



These intermolecular forces as a group are referred to as van der Waals forces.


van der Waals Forces

• Dipole-dipole interactions

• Hydrogen bonding

• London dispersion forces


Ion-Dipole Interactions

• Ion-dipole interactions (a fourth type of force), are important in solutions of ions.

• The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents.


Dipole-Dipole Interactions

• 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.


Dipole-Dipole Interactions

The more polar the molecule, the higher is its boiling point.


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), it does happen that they occasionally wind up on the same side of the atom.



At that instant, then, the helium atom is polar, with an excess of electrons on the left side and a shortage on the right side.



Another helium nearby, then, would 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.



London dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and an induced dipole.



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

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


Factors Affecting London Forces

• 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.


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.


Which Have a Greater Effect?Dipole-Dipole Interactions or Dispersion Forces

• If two molecules are of comparable size and shape, dipole-dipole interactions will likely the dominating force.

• If one molecule is much larger than another, dispersion forces will likely determine its physical properties.


How Do We Explain This?

• The nonpolar series (SnH4 to CH4) follow the expected trend.

• The polar series follows the trend from H2Te through H2S, but water is quite an anomaly.

IntermolecularForces

Comparing IMF

1. The dipole moments of acetonitrile, CH3CN, and methyl iodide, CH3I, are 3.9 D and 1.62 D, respectively. (a)Which of these substances has greater dipole–dipole attractions among its molecules? (b)Which of these substances has greater London dispersion attractions? (c)The boiling points of CH3CN and CH3I are 354.8 K and 315.6 K, respectively. Which substance has the greater overall attractive forces?

2. Of Br2, Ne, HCl, HBr, and N2, which is likely to have(a)the largest intermolecular dispersion forces

(b) the largest dipole–dipole attractive forces?



Hydrogen Bonding

• The dipole-dipole interactions experienced when H is bonded to N, O, or F are unusually strong.

• We call these interactions hydrogen bonds.


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 hydrogen nucleus is exposed.


Summarizing Intermolecular Forces

Hydrogen Bonding

1. In which of the following substances is hydrogen bonding likely to play an important role in determining physical properties: methane (CH4), hydrazine (H2NNH2), methyl fluoride (CH3F), or hydrogen sulfide (H2S)?

2. In which of the following substances is significant hydrogen bonding possible: methylene chloride (CH2Cl2), phosphine (PH3), hydrogen peroxide (HOOH), or acetone (CH3COCH3)?


Predicting the Types and Relative Strengths of Intermolecular Attractions

1. List the substances BaCl2, H2, CO, HF, and Ne in order of increasing boiling points.

2. Identify the intermolecular attractions present in the following substances, and select the substance with the highest boiling point: CH3CH3, CH3OH, and CH3CH2OH.



Intermolecular Forces Affect Many Physical Properties

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


Viscosity• 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.


Surface Tension

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


Phase Changes


Energy Changes Associated with Changes of State

The heat of fusion is the energy required to change a solid at its melting point to a liquid.



The heat of vaporization is defined as the energy required to change a liquid at its boiling point to a gas.



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

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

Calculating ΔH for Temperature and Phase Changes

1. Calculate the enthalpy change upon converting 1.00 mol of ice at –25 oC to water vapor (steam) at 125 oC under a constant pressure of 1 atm. The specific heats of ice, water, and steam are 2.03 J/g-K, 4.18 J/g-K, and 1.84 J/g-K, respectively. For H2O, ΔHfus = 6.01 kJ/mol and ΔHvap = 40.67 kJ/mol.

2. What is the enthalpy change during the process in which 100.0 g of water at 50.0 oC is cooled to ice at –30.0 °C ?



Vapor 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.


Vapor Pressure

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


Vapor Pressure

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


Vapor Pressure

• The boiling point of a liquid is the temperature at which it’s vapor pressure equals atmospheric pressure.

• The normal boiling point is the temperature at which its vapor pressure is 760 torr.

1. Use Figure 11.24 to estimate the boiling point of diethyl ether under an external pressure of 0.80 atm.

2. At what external pressure will ethanol have a boiling point of 60 °C?



Phase Diagrams

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


Phase Diagrams

• The circled line is the liquid-vapor interface.• It starts at the triple point (T), the point at

which all three states are in equilibrium.


Phase Diagrams

It ends at the critical point (C); above this critical temperature and critical pressure the liquid and vapor are indistinguishable from each other.


Phase Diagrams

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


Phase Diagrams

• The circled line in the diagram below is the interface between liquid and solid.

• The melting point at each pressure can be found along this line.


Phase Diagrams• Below the triple point the substance cannot

exist in the liquid state.• Along the circled line the solid and gas

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


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.


Phase Diagram of Water

• The slope of the solid-liquid line 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

Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm; CO2 sublimes at normal pressures.


Phase Diagram of Carbon Dioxide

The low critical temperature and critical pressure for CO2 make supercritical CO2 a good solvent for extracting nonpolar substances (like caffeine)


Solids

• We can think of solids as falling into two groups:– crystalline, in which

particles are in highly ordered arrangement.


Solids

• We can think of solids as falling into two groups:– amorphous, in which

there is no particular order in the arrangement of particles.


Attractions in Ionic CrystalsIn ionic crystals, ions pack themselves so as to maximize the attractions and minimize repulsions between the ions.


Types of Bonding in Crystalline Solids


Covalent-Network andMolecular 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.


Covalent-Network andMolecular Solids

• Graphite is an example of a molecular solid, in which atoms are held together with van der Waals forces.– They tend to be softer and have lower melting

points.


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 delocalized throughout the solid.

Putting Concepts Together

1. The substance CS2 has a melting point of –110.8 C and a boiling point of 46.3 °C. Its density at 20 °C is 1.26 g/cm3. It is highly flammable.

• (a) What is the name of this compound?

• (b) List the intermolecular forces that CS2 molecules would have with each other.

• (c) Predict what type of crystalline solid CS2(s) would form.

• (d) Write a balanced equation for the combustion of this compound in air. (You will have to decide on the most likely oxidation products.)

• (e) The critical temperature and pressure for CS2 are 552 K and 78 atm, respectively. Compare these values with those for CO2 (Table 11.5), and discuss the possible origins of the differences.

• (f) Would you expect the density of CS2 at 40 °C to be greater or less than at 20 °C? What accounts for the difference?


Ans.

(a) The compound is named carbon disulfide, in analogy with the naming of other binary molecular compounds such as carbon dioxide . (b) Only London dispersion forces affect CS2; it does not have a dipole moment, based upon its molecular shape, and obviously cannot undergo hydrogen bonding.(c) Because CS2(s) consists of individual CS2 molecules, it will be a molecular solid.(d) The most likely products of the combustion will be CO2 and SO2. (Sections 3.2 and 7.8) Under some conditions SO3 might be formed, but this would be the less likely outcome. Thus, we have the following equation for combustion:

CS2(l) + 3 O2(g) → CO2(g) + 2 SO2(g)

(e) The critical temperature and pressure of CS2 (552 K and 78 atm) are both higher than those given for CO2 in Table 11.5 (304 K and 73 atm). The difference in critical temperatures is especially notable. The higher values for CS 2 arise from the greater London dispersion attractions between the CS2 molecules compared with CO2. These greater attractions are due to the larger size of the sulfur compared to oxygen and therefore its greater polarizability.(f) The density would be lower at the higher temperature. Density decreases with increasing temperature because the molecules possess higher kinetic energies. Their more energetic movements result in larger average distances between molecules, which translate into lower densities.


© 2009, prentice-hall, inc. chapter 11 intermolecular forces, liquids, and solids

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