Chapter 11. Intermolecular Forces, Liquids, and Solids · Chapter 11. Intermolecular Forces, Liquids, and Solids ... much weaker than the forces of attraction that hold one water

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Chapter 11. Intermolecular Forces, Liquids, and Solids

NOTE – THIS CHAPTER IS ONE OF THE TOP TOPICS ON AP EXAM!!! NOT ONLY DO YOU NEED TO FOCUS ON THEORY (NO MATH IN THIS TOPIC) BUT YOU MUST READ THE FIGURES TOO!!! 11.1 A Molecular Comparison of Gases, Liquids and Solids SG Questions – Read p. 442-446. Answer the following questions

1.Define intermolecular forces and intramolecular. What is a BIG difference between them?

2. What determines the state of the substance? 3. Fill in the chart to Summarize Properties of States of Matter

State of

Matter

Shape &

Volume

Expands Compressible Flows Diffusion

Solid

Liquid

Gas

4. List from least to greatest of a solid, liquid and gas on the strength of interparticle attractions

5. What happens to the kinetic energy of particles of a gas when temperature is decreased?

6. How does pressure play to affect when the pressure of a gas is increased?

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Physical properties of Gases, Liquids, Solids: • Gases are highly compressible and assume the shape and volume of their containers. • Gas molecules are far apart and do not interact much with one another.

• Liquids are incompressible; they assume the shape, but not the volume of the container. • Liquid molecules are held together more closely than gas molecules, the molecules

slide past each other. • Solids are incompressible and have a definite shape and volume. • Solid molecules are packed closely together. • The molecules are so rigidly packed that they cannot slide past each other. • Solids with highly ordered structures are said to be crystalline.

• Converting a gas into a liquid or solid requires the molecules to get closer to each other. • We can accomplish this by cooling or compressing the gas. • Converting a solid into a liquid or gas requires the molecules to move further apart. • We can accomplish this by heating or reducing the pressure on the solid.

Now it is time to consider the forces that condense matter. The forces that hold one molecule to

another molecule are referred to as intermolecular forces (IMFs). These forces arise from

unequal distribution of the electrons in the molecule and the electrostatic attraction between

oppositely charged portions of molecules. We briefly visited the IMFs earlier when discussing

the nonideal behavior of gases. These forces cause changes of state by causing changes among

the molecules NOT within them.

Physical properties such as melting points, boiling points, vapor pressures, etc. can be attributed

to the strength of the intermolecular attractions present between molecules.

It works like this: the lower the boiling point (or vapor pressure or melting point), the

weaker the intermolecular attractions; the higher the boiling point, the stronger the

intermolecular attractions. For example, gasoline evaporates much more quickly than water.

Therefore, the intermolecular attractive forces that hold one gasoline molecule to another are

much weaker than the forces of attraction that hold one water molecule to another water

molecule. In fact, water molecules are held together by the strongest of the intermolecular

attractive forces, hydrogen bonds. Hydrogen bonds are not true bonds—they are just forces of

attraction that exist between a hydrogen atom on one molecule and the unshared electron pair on

fluorine, oxygen or nitrogen atoms of a neighboring molecule.

Intermolecular Forces, Liquids, and Solids 3

11.2 Intermolecular Forces (IMF) – VERY IMPORTANT!!! SG Questions – Read p. 446-455. Answer the following questions

1. Which are stronger intermolecular forces or intramolecular forces?

2. What is the relationship between boiling point, melting point and intermolecular forces

3. What are the 3 IMF and who were they founded by?

4. What is London Dispersion (LD)?

5. What is polarizability? How does polarizability affect this force?

6. What type of relationship does the strength of LD have with atomic/molecular size?

7. Define what dipole-dipole forces are.

8. How do you determine the strength of intermolecular attractions when the molecules

have approximately equal mass and size?

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9. What is hydrogen bonding? FOR WHAT ELEMENTS IS IT ONLY CONSIDERED

HYDROGEN BONDING?

10. List the intermolecular forces (HB,DD,LD) by increasing strength.

11. What is the importance of HB in other sciences?

12. When do ion-dipole forces exists? See Figure 11.13

13. What happens to the strength of the IMF if molecule polarity increases?

14. If two molecules have the same forces what you use to determine which is stronger/

Figure 11.3 is a great explanation and visual of intramolecular forces

• Intermolecular forces are much weaker than ionic or covalent bonds. • When a substance melts or boils, intermolecular forces are broken. • When a substances condenses, intermolecular forces are formed. • A high boiling point indicates strong attractive forces. • A high melting point indicates strong attractive forces. • Van der Waals forces (3 TYPES) are the intermolecular forces that exist between neutral molecules.

1. London-dispersion forces 2. Dipole-dipole forces 3. Hydrogen-bonding forces.


London-Dispersion Forces

• These are the weakest of all intermolecular forces. • London-dispersion forces exist between ALL MOLECULES!!!! • What affects the strength of a dispersion force? • Molecules must be very close together for these attractive forces to occur. • Polarizability is the ease with which an electron distribution can be deformed. • London-Dispersion Forces are all NONPOLAR molecules!!! The larger the molecule (the greater the number of electrons) the more polarizable it is. • London-dispersion forces increase as molecular weight increases. • London-dispersion forces depend on the shape of the molecule. Nonpolar molecules have no natural attraction for each other. • To better understand London-Dispersion Forces, lets examine the halogens. Halogens exist

as diatomic molecules at room temperature and atmospheric conditions. F2 and Cl2 are gases,

Br2 is a liquid and I2 is a solid. Why?

• All of these molecules are completely nonpolar and according to theory, not

attracted to each other, so one might predict they would all be gases at room

temperature.

• Bromine exists as a liquid at room temperature simply because there is a greater

attractive force between its molecules than between those of fluorine or chlorine.

Why?

• Bromine is larger than fluorine or chlorine; it has more electrons and is thus more

polarizable. Electrons are in constant motion so it is reasonable that they may

occasionally “pile up” on one side of the molecule making a temporary negative pole

on that end, leaving a temporary positive pole on the other end.

• Iodine is a solid since it is larger still, has even more electrons, is thus even more

polarizable and the attractive forces are thus even greater.

Every C-H molecules have LD!!!!!

Dipole-Dipole Forces

• Dipole-dipole forces exist between neutral POLAR molecules. To determine if a a molecule is POLAR, you need to know their electronegativity (ability to gain electrons) numbers. Dipole-dipole forces have attraction that enables two polar molecules to attract one another.

Polar molecules are those which have an uneven charge distribution since their dipole moments

do not cancel. Compounds exhibiting this type of IMF have higher melting and boiling points

than those exhibiting weaker IMFs.

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• Polar molecules attract each other. “like dissolves like” • The partially positive end of one molecule attracts the partially negative end of another. • Polar molecules need to be close together to form strong dipole-dipole interactions.

Think Oil vs. Water – Do they mix? Why? • If two molecules have about the same mass and size, then dipole-dipole forces increase with

increasing polarity. (See figure below) • For molecules of similar polarity, those with smaller volumes often have greater dipole-

dipole attractions.

For example: Hydrochloric acid molecules are held to each other by this type of force. HCl—the

chlorine pulls the electrons in the bond with greater force than hydrogen so the molecule is polar

in terms of electron distribution. Two neighboring HCl molecules will align their oppositely

charged ends and attract one another.

Figure 11.8- p.449 Moving left to right, do the dispersion forces get stronger, weaker, or

roughly stay the same? (See answer p. A-39 in back of book)

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Hydrogen Bonding

Hydrogen bonding—the force of attraction between the hydrogen atom of one molecule and an unshared electron pair on F, O, or N of a neighboring molecule (a special case of dipole-dipole). This is the strongest IMF.

Experiments show that the boiling points of compounds with H–F, H–O, and H–N bonds

are abnormally high.

Never confuse hydrogen bonding with a bonded hydrogen. The unique physical properties of

water are due to the fact that it exhibits hydrogen bonding. As a result of these attractions,

water has a high boiling point, high specific heat, and many other unusual properties.

• BP increases with increasing molecular mass [since the more electrons in a molecule, the

more polarizable the cloud {more about that in the next section}, the stronger the IMFs,

therefore the more Energy needed to overcome those increased attractions and vaporize, thus

the higher the boiling point. That’s how it is supposed to work!]. Hydrogen bonding, that’s

why!

Figure 11.10- p.450 To form a Hydrogen - bond

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Ion-Dipole Forces

• An ion-dipole force is an interaction between an ion (e.g., Na+) and the partial charge on the

end of a polar molecule/dipole (e.g., water). DISSOCIATION!!!!

• Example: NaCl (aq) Draw Figure 11.13 – p.452

Comparing Intermolecular Forces

Dispersion forces are found in all substances.

• Their strength depends on molecular shapes and molecular weights. Dipole-dipole forces add to the effect of dispersion forces. • They are found only in polar molecules. H-bonding is a special case of dipole-dipole interactions. • It is the strongest of the intermolecular forces involving neutral species. • H-bonding is most important for H compounds of N, O, and F. If ions are involved, ion-dipole (if a dipole is present) and ionic bonding are possible.

• Ion-dipole interactions are stronger than H-bonds.. Keep in mind: IONIC bonds are stronger forces than van der waals forces (London, Dipole-Dipole, H-Bonding)!


11.3 Some Properties of Liquids SG Questions – Read p. 455-457. Answer the following questions

1. Define viscosity

2. How does temperature affect the viscosity?

3. What is the relationship between the flowing of the liquid with greater liquid viscosity?

4. Why do water droplets assume an almost spherical shape?

5. Define surface tension

6. Why is the surface tension of mercury greater than the surface tension of water?

7. What are the differences between cohesive and adhesive forces?

8. Why does water have meniscus have a U-shape but mercury has an inverted U-shape,

how are the forces different for mercury and water?

9. Why does water rise in a tube?

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Viscosity

• Viscosity is the resistance of a liquid to flow. (Maple Syrup or Honey) • A liquid flows by sliding molecules over one another. Resistance to flow (molecules with strong intermolecular forces). Increases with molecular

complexity [long C chains (sugars & alcohols) get tangled and larger electron clouds are more

polarizable due to the presence of additional electrons] and increased with increasing IMFs. • Viscosity depends on the attractive forces between molecules. • The stronger the intermolecular forces are, the higher the viscosity. • Viscosity usually decreases with an increase in temperature.

Surface Tension

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

The resistance to an increase in its surface area (polar molecules). High ST indicates strong

IMFs. Molecules are attracted to each OTHER. A molecule in the interior of a liquid is attracted

by the molecules surrounding it, whereas a molecule at the surface of a liquid is attracted only by

the molecules below it and on each side. • Stronger intermolecular forces cause higher surface tension. • Water has a high surface tension (H-bonding). • Hg(l) has an even higher surface tension (there are very strong metallic bonds between Hg atoms). Cohesive and adhesive forces: • Cohesive forces are intermolecular forces that bind molecules to one another. • Adhesive forces are intermolecular forces that bind molecules to a surface. (water bent towards the side of container) • Illustrate this by looking at the meniscus in a tube filled with liquid. • The meniscus is the shape of the liquid surface. • If adhesive forces are greater than cohesive forces, the meniscus is U-shaped (e.g.,

water in glass). • If cohesive forces are greater than adhesive forces, the meniscus is curved downwards

(e.g., Hg(l) in glass).


Figure 11.19 – Meniscus shapes for water and mercury

11.4 Phase Changes - Very Important

SG Questions – Read p. 457-461. Answer the following questions

1. What are the 6 phases changes and indicate if they are endothermic or exothermic (3

endo and 3 exo)

2. By increasing the temperature of a solid how does it affect the particles what changes

occur?

3. What is fusion?

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4. How are the freedom of motion measured from different phases changes?

5. What is heat of vaporization and what are the units that can be used for this calculation?

6. When can the heat of sublimation be used, what phases does it change into?

7. What is the heating curve?

8. What us the difference between specific heat and specific heat capacity?

9. List of examples of supercooled and superheated

10. Define critical pressure and critical temperature

11. What is the relationship between intermolecular forces, critical temperature, and

molecular weight?

12. What is the importance of a supercritical fluid?


• Phase changes are changes of state. • Matter in one state is converted into another state. Phase Changes Energy Changes • sublimation: solid → gas. (ENDO)

• melting: solid → liquid. (ENDO) • vaporization: liquid →gas. (ENDO) • deposition: gas →solid. (EXO) • condensation: gas → liquid. (EXO) • freezing: liquid → solid. (EXO)

***When you cant remember these, THINK OF HOW WATER CHANGES***

ENDOTHERMIC – needing energy EXOTHERMIC – releasing energy It takes MORE energy to completely separate (break) molecules, than to combining (form) them . Figure 11.20 – Phase Changes

Volatile substances (liquid take evaporate easily - changes to a gas easily)— have high VP

(vapor pressure), thus weak IMFs. These liquids evaporate readily from open containers since

they have so little attraction for each other. It takes very little energy being absorbed in order for

them to escape the surface of the liquid. The heat energy absorbed from a warm room is usually

enough to make these substances evaporate quickly. If there is an odor to the substance, these

are the liquids you smell almost as soon as you open the bottle! The molecules have been

banging against the lid wanting out!

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• VP increases significantly with temperature! Heat ‘em up, speed ‘em up, move ‘em

out! Increasing the temperature increases the KE which facilitates escape AND the speed of the

escapees! They bang into the sides of the container with more frequency [more of them escaped] The higher the temperature, the higher the average kinetic energy, the faster the liquid evaporates. • As MM ↑ VP ↓ ……BECAUSE as molecules increase in molar mass, they also increase

in the number of electrons. As the number of electrons increase, the polarizability of the

molecule increases so more dispersion forces exist, causing stronger attractions to form between

molecules. This decreases the number of molecules that escape and thus lowers the VP. Heating Curves – See Figure 11.22 p. 459

• Plot of temperature change versus heat added over a period of time.

• During a phase change adding heat causes no temperature change. • The added energy is used to break intermolecular bonds rather than cause a

temperature change.

FREAKS—changes of state do not always form at the exact MP and BP’s.

• supercooled—oxymoron—the substance is at a temperature below its FP, yet it remains a

liquid. Usually happens when the cooling has been gradual and the degree of organization

needed to form crystals hasn’t happened. At some point, organization happens and the solid

is readily formed, the temperature rises back to the MP as the heat of crystallization is

released.

superheated—another oxymoron—the substance is at a temperature above its BP, yet it

remains a liquid. Usually happens when heated very rapidly [microwave oven] and bubbles

form in the interior with high internal pressures. They often burst before reaching the

surface making quite a mess of things! Really ruins an experiment.


11.5 Vapor Pressure – VERY IMPORTANT!!! SG Questions – Read p. 461-464. Answer the following questions

1. When is vapor pressure attained?

2. Figure 11.23 – Summarize what is occurring!!

3. What is dynamic equilibrium?

4. What is the vapor pressure of a liquid?

5. In what certain container does vaporization occur in?

6. If liquids are known to evaporate readily are known as _________________.

7. Why does hot water evaporate faster than cold water? (KE)

8. Figure 11.24 – Effect of T on the kinetic energies in liquid.

As the temperature increases, does the rate of molecules escaping into gas phase increase or

decrease? (Read p.462 and see see answer on page A-40 in back of book)

Label what the blue and red areas represent!

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9. What are the specific requirements for normal boiling point?

10. How does pressure affect the boiling point of a liquid or gas?

11. See Figure 11.25 p. 463 – Summarize how vapor pressure with temperature for 4

different substances differs greatly in volatility


Explaining Vapor Pressure on the Molecular Level (Figure 11.23)

• Dynamic equilibrium - a condition in which two opposing processes occur simultaneously at equal rates.

• In this case, it is the point when as many molecules escape the surface as they strike the surface. • Vapor pressure of a liquid is the pressure exerted by its vapor when the liquid and vapor

are in dynamic equilibrium. • The pressure of the vapor at this point is called the equilibrium vapor pressure.

Vapor Pressure and Boiling Point (Figures 11.24-11.25)

• Liquids boil when the external pressure at the liquid surface equals the vapor pressure. • The normal boiling point is the boiling point at 760 mm Hg (1 atm). • The temperature of the boiling point increases as the external pressure increases.

• Two ways to get a liquid to boil are to: • increase temperature or decrease pressure. (Remember: this is the conditions of ideal gases!)

11.6 Phase Diagrams – VERY IMPORTANT!!!

SG Questions – Read p. 442-446. Answer the following questions

1. Define a phase diagram

2. What does the phase diagram tell you?

3. What is the significance of the critical point what does it determine?

4. What does the sublimation curve separate?

5. What does the melting phase separate?

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6. How does one know that they can identify the normal boiling point and normal melting

point?

7. What is the triple point?

Phase Diagrams—closed systems A phase diagram is a plot of pressure vs. temperature summarizing all equilibrium between phases.

• Phase diagrams tell us which phase will exist at a given temperature and pressure. • Features of a phase diagram include: • vapor-pressure curve: generally as temperature increases, vapor pressure increases. • normal melting point—the temperature at which the VP solid = VP liquid (1atm)

• normal boiling point—the temperature at which the VP liquid = VP gas (1 atm) • triple point: temperature and pressure at which all three phases are in equilibrium.

Figure 11.27 – Phase Diagram for a Substance


Critical Temperature and Pressure

Gases may be liquefied by increasing the pressure at a suitable temperature.

• Critical temperature is the highest temperature at which a substance can exist as a liquid. • Critical pressure is the pressure required for liquefaction at this critical temperature. • The greater the intermolecular forces, the easier it is to liquefy a substance. • Thus, the higher the critical temperature.

Critical point: critical temperature and pressure coordinates (for water, Tc = 374°C and 218

atm).

Phase Diagrams of H2O and CO2

Figure 11.28 – Phase Diagram for Water Let’s practice Drawing

Water:

• In general, an increase in pressure favors the more compact phase of the material. • This is usually the solid. • Water is one of the few substances whose solid form is less dense than the liquid form. • The melting point curve for water slopes to the left. • The triple point occurs at 0.0098◦C and 4.58 mm Hg. • The normal melting (freezing) point is 0◦C. • The normal boiling point is 100◦C. • The critical point is 374◦C and 218 atm. Water is a freak! The solid-liquid line tilts to the left [negative slope] since its solid is less

dense than its liquid phase—ice floats. (Usually, most solids sink as they are more dense

than liquids).

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Figure 11.29 p. 466 – Phase Diagram for CO2

Most substances have a solid-liquid line that has a positive slope since their

solid phase is more dense than the liquid. This one is for carbon dioxide.

EACH PHASE BOUNDARY REPRESENTS AN EQUILIBRIUM SET OF PRESSURE

AND TEMPERATURE CONDITIONS!! Be sure and use the word equilibrium in your

writing as you explain phase diagrams.


12.1 Classifications of Solids • Solids are classified by the types of bonds that hold the atoms in place. • Metallic solids: held together by a delocalized “sea” of shared valence electrons.

• Ionic solids: held together by mutual attraction between cations and anions. • Covalent-network solids: held together by an extended network of covalent bonds. • Molecular solids: held together by relatively weak intermolecular forces. • Some solids are harder to categorize: • Polymers: long chains of atoms held together by covalent bonds. • The chains are often held to one another by weaker intermolecular forces.

• Nanomaterials: solids whose crystals have dimensions on the order of 1–100nm.

12.2 Structures of Solids Crystalline and Amorphous Solids

• Solids may be classified according to their level of order. • Crystalline solids have atoms and ions arranged in an orderly repeating pattern.

• Examples: NaCl, quartz, diamond • Amorphous solids lack the order of crystalline solids. • Examples: rubber and glass Unit Cells and Crystal Lattices

• Crystalline solids have an ordered, repeating structure. • The smallest repeating unit in a crystal is a unit cell.

• The unit cell is the smallest unit with all the symmetry of the entire crystal. • The three-dimensional stacking of unit cells is the crystal lattice.

Filling the Unit Cell

• A crystal structure is defined by the lattice as well as the location of the atoms relative to the lattice

points. • Simplest case: a crystal structure made of identical atoms with each atom exactly coincident with

a lattice point. • Here the crystal structure and the lattice points have the same pattern. • Many metallic elements have this structure. • Only elements can form this structure.

12.3 Metallic Solids • Metallic solids or metals consist solely of metal atoms. • Metallic bonding results from delocalization of valence electrons throughout the solid. • The metal nuclei are seen to exist in a “sea” of delocalized valence electrons. • Important physical properties of pure metals: • malleable: can be hammered into thin sheets

• ductile: can be pulled into wires • good electrical and thermal conductivity

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The Structures of Metallic Solids

Close Packing

• The sharing of valence electrons makes it favorable for the atoms in a metal to pack closely together.

• Layers of atoms are efficiently packed by surrounding each atom by six neighbors. • A three-dimensional structure is formed by stacking additional layers on top of the base layer. • The second layer packs most efficiently if the atoms sit in the depressions formed by the

atoms in the first layer.

Alloys

• Alloys contain more than one element and have the characteristic properties of metals. • Pure metals and alloys have different physical properties.

• An alloy of gold and copper is used in jewelry (the alloy is harder than the relatively soft pure 24 karat gold).

• 14 karat gold is an alloy containing 58% gold. • Alloys can be divided into 4 categories:

• Substitutional alloys (the solute atoms take the positions normally occupied by a solvent atom) • The atoms must have similar atomic radii. • The elements must have similar bonding characteristics. • Interstitial alloys (the solute occupies interstitial sites in the metallic lattice) • One element (usually a nonmetal) must have a significantly smaller radius than the other (in

order to fit into the interstitial site).

• The alloy is much harder, stronger and less ductile than the pure metal (increased bonding between nonmetal and metal).

• An example is steel (contains up to 3% carbon). • Other elements may also be added to make alloy steels. • Addition of V and Cr increases the strength of the steel and improves its resistance to

stress and corrosion.

• The most important iron alloy is stainless steel. It contains C, Cr (from ferrochrome, FeCr2), and Ni.

• Heterogeneous alloys: The components are not dispersed uniformly (e.g., pearlite steel has two

phases: almost pure Fe and cementite, Fe3C). • Intermetallic compounds: homogeneous alloys with definite properties and compositions.

• Examples include: • Ni3Al (a major component of jet aircraft engines). • Cr3Pt (used to coat razor blades to increase hardness and ability to maintain a sharp edge), • Co5Sm (used in permanent magnets in lightweight headsets). • LaN5 (used as the anode in nickel-metal hydride batteries).

12.4 Metallic Bonding

Electron-Sea Model

• A simple model that accounts for many properties of metals is the electron-sea model. • An array of metal cations exist in a “sea” of valence electrons. • No electrons are localized between any two metal atoms.


• Therefore, the electrons can flow freely through the metal. • Without any definite bonds, the metals are easy to deform (and are malleable and ductile) and

exhibit high thermal and electrical conductivity.

Molecular-Orbital Model1,2

• Problems with the electron-sea model: • As the number of electrons increases, the strength of bonding should increase, and the melting

point should increase. • However, group 6B metals (at the center of the transition metals) have the highest melting points

in their respective periods.

• Consider how the molecular-orbital diagram for a chain of lithium atoms changes as we increase the length of the chain. • As the chain length increases, the number of molecular orbitals increases. • If the chain becomes very long, there are so many molecular orbitals that the energy separation

between orbitals becomes vanishingly small.

• As the number of orbitals increases, their energy spacing decreases. • As the chain length goes to infinity the allowed energy states become a continuous band. • The complex electronic structure of a bulk solid is call a band structure. • Many of the properties of metals can be understood in terms of their band structure.

• The available electrons do not completely fill the band of orbitals. • Therefore, electrons can be promoted to unoccupied energy bands.

• Because the energy differences between orbitals are small, the promotion of electrons requires little energy.

12.5 Ionic Solids

• Ionic solids are held together by ionic bonds: electrostatic attraction between cations and anions. • The high melting and boiling points reflect the strength of ionic bonds.

• The strength depends on the charges and sizes of the ions. • Attractions increase in strength as the charges of the ions go up. • Attractions also increase as the ions get smaller. • Valence electrons in ionic compounds are confined to the anions rather than being delocalized. • Thus, ionic solids tend to be electrical insulators. • Ionic solids tend to be brittle.

• Stress applied to an ionic solid may shift the alignment of ions and create repulsive interactions between ions of like charge.

Structures of Ionic Solids

• Ionic solids tend to have close-packed arrangements of ions.

• The arrangement is different from that of metallic solids due to the different radii and different charges of the cations and anions.

1 Figure 12.21

2 Figure 12.23

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12.6 Molecular Solids • Molecular solids consist of atoms or molecules held together by intermolecular forces. • Weak intermolecular forces give rise to low melting points. • Intermolecular forces include dipole-dipole, dispersion forces and H-bonds.

• Molecular solids are usually soft. • They are often gases or liquids at room temperature. • Efficient packing of molecules is important (since they are not regular spheres). • Molecular solids show poor thermal and electrical conductivity. • Examples: Ar(s), CH4(s), CO2(s), sucrose.

12.7 Covalent-Network Solids

• Covalent-network solids consist of atoms held together, in large networks or chains, with covalent bonds.

• They have much higher melting points and are much harder than molecular solids. • This is a consequence of the strong covalent bonds that connect the atoms. • Examples are diamond, graphite, quartz (SiO2), and silicon carbide (SiC).

• In diamond:

• each C atom is tetrahedral. • Diamond is hard, and has a high melting point (3550 °C). • In graphite: • each C atom is arranged in a planar hexagonal ring. • layers of interconnected rings are placed on top of each other. • electrons move in delocalized orbitals (good conductor).

• the layers are held together by weak dispersion forces. • They slide easily past each other. • Graphite is a good lubricant.

Semiconductors

• Certain solids conduct electricity but not as well as metals; these are called semiconductors. • Semiconductors can be divided into two classes.

• Elemental semiconductors (made of only one type of atom) • Elemental semiconductors include silicon, germanium, and gray tin. • These elements adopt the crystal structure of diamond.

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Chapter 11. Intermolecular Forces, Liquids, and Solids · Chapter 11. Intermolecular Forces, Liquids, and Solids ... much weaker than the forces of attraction that hold one water