Chapter 10 - LIQUIDS and SOLIDS
Learning Goals:
 Learn about intermolecular forces – dipole-dipole, London dispersion, and H-bonds;
 Relationships between intermolecular forces and vapor pressure, boiling point, enthalpy of vaporization, surface tension, capillary action, and viscosity of liquids;
 Vapor pressure-Temperature Equilibria and Phase diagrams;
 Structure and bonding in solids;
 Types of crystalline solids and unit cells;
10.1 Intermolecular Forces
Intermolecular forces are attractive forces between molecules (or particles) that make up a substance and they are especially important in liquids and solids. Attractive intermolecular forces enable real gases to condense and form liquids. Because molecules or particles in gases are generally very far apart from one another, the influence of intermolecular forces is negligible and is normally ignored. However, because of the close proximity of molecules to each other in the liquids and solids, intermolecular forces play important roles in determining the physical properties of liquids and solids.
Intermolecular forces are responsible for many properties of liquids, such as surface tension, viscosity, capillary action, vapor pressure, boiling point, and enthalpy of vaporization.
Types of Intermolecular Forces in Liquids:
 Dispersion (London) forces;
 Permanent dipole-dipole interactions;
 Hydrogen bonds;
Electron cloud in molecules can be dispersed and polarized, regardless of the dipole characteristics of the molecules. Thus, nonpolar molecules may exhibit dipole-dipole attractions through charge dispersion processes. Intermolecular forces that result from charge dispersion is called dispersion forces. For example, a nonpolar molecule like Br
2
acquires an instantaneous (non-permanent) dipole moment as a result of charge dispersion and molecular polarization. As one molecule becomes polarized, it induces the neighboring molecules to become polarized, resulting in dipole-dipole attractions. Dispersion (London) forces occur in all substances, but in substances containing nonpolar molecules, they represent the only type of intermolecular attractions.
The strength of dispersion forces depends on the polarizability of the molecules. Larger and elongated molecules tend to be more easily polarized because the valence electrons are less firmly held by the nuclei. Therefore, substances containing large molecules are expected to exhibit stronger dispersion forces and higher boiling points. For example, dispersion forces and boiling points increase down groups, such as;
He < Ne < Ar < Kr < Xe; F
2
< Cl
2
< Br
2
< I
2
, and CH
4
< SiH
4
< GeH
4
< SnH
4
1

Dispersion forces in hydrocarbons increase as molecular masses increase, and their boiling points also increase accordingly:
(B.p.,
CH
4
< C
2
H
6
< C o C): -162 -89
4
H
10 6
H < C
8
H
18
< C
10
H
22
methane ethane butane hexane octane decane
-1
< C
14
69 126 174
Molecular shape also influences the strength of dispersion forces. For example, npentane and neopentane (2,2-dimethylpropane) both have the molecular formula C
5
H
12
and are nonpolar, but n-pentane has an elongated shape, whereas neopentane is approximately spherical. n-Pentane has a boiling point of 36 o C compared to 10 o C for neopentane, which implies stronger dispersion forces in n-pentane due to its elongated molecular shape that makes it more readily polarizable.
CH
3
CH
3
CH
2
CH
2
CH
2
CH
3

H
3
C—C—CH
3
 n-pentane
(elongated; b.p. = 36 o C)
CH
3
neopentane (2,2-dimethylpropane)
(spherical; b.p. ~ 10 o C).
As molecules become larger, the contribution of London dispersion forces increases.
———————————————————————————————————————
Substances MM  (D) Dispersion % Permanent  H v
B.Pt.
Forces (%) dipole-dipole kJ/mol) (K)
———————————————————————————————————————
F
2
HCl
HBr
38.00
36.46
80.92
0
1.08
0.82
100
81.4
94.5
0
18.6
5.5
6.86
16.15
17.61
85.01
188.11
206.43
HI 127.91 0.44 99.5 0.5 19.77 237.80
———————————————————————————————————————
Contributions of permanent dipole-dipole interactions: HCl > HBr > HI > F
2
.
Contributions of dispersion London forces: F
2
> HI > HBr > HCl
Molecules with polar bonds and asymmetric structures often have permanent dipole
moments – they have positive and negative ends. Such molecules are attracted to each other electrostatically by
. In the liquid state, polar molecules orient themselves in such a way that attractions between oppositely charged ends of the molecules are maximized, while repulsions between like charged ends are minimized. Permanent dipoledipole attractions are important intermolecular forces in substances containing small molecules where the dipole moments are relatively large. Dipole moment decreases with increasing molecular size, such as the following: HF > HCl > HBr > HI.
In substances containing small polar molecules, permanent dipole-dipole attractions play a major influence on liquid properties. For example, HCl has a higher boiling point and enthalpy of vaporization than F
2
due to the former having a permanent dipole-dipole forces, in addition to the dispersion (London) forces. Acetone, (CH polar molecules, has a higher boiling point (56.2 o
3
)
2
CO, an organic solvent containing
C) than butane (C
4
H
10
, b.p. ~-1 o C), which contains nonpolar molecules. Both substances have the same molecular mass and they are
2

expected to exhibit similar dispersion forces. However, the overall intermolecular attractions between acetone molecules are stronger than between butane molecules due to the permanent dipole-dipole attractions in acetone.
CH
3
CH
2
CH
2
CH
3
n-butane
H
3
C
C — O (acetone,  > 0; b.p. 56.2 o C)
(  = 0, b.p. -1 o C) H
3
C
occurs in molecules containing N—H, O—H, or H—F bonds. These are relatively strong dipole-dipole interactions, with energy in the range of 15 - 40 kJ/mol.
Substances containing hydrogen bonds usually occur as liquids (under normal conditions) with relatively high boiling points, enthalpy of vaporization, surface tension, and viscosity. For example, H
2
O, NH
3
and HF, though contain the smallest molecules, they exhibit the highest boiling points in their respective groups. Hydrogen bonds are responsible for many special properties of water, such as high boiling point, enthalpy of vaporization, specific heat, surface tension, and capillary action. We regulate our body temperature through sweating, which removes excess body heat through evaporation.
Hydrogen bonds are important intra- and inter-molecular forces in biological macromolecules. Collectively, they are responsible for the native and functional structures of proteins, DNA, RNA, as well as carbohydrates and celluloses.
Exercise-1:
1. Indicate the type of intermolecular forces the following molecules will form with each other in the liquid or solid state.
(a) CH
(f) CH
2
4
O
(b) CH
(g) NH
3
2
OH
OH
(c) CH
3
F
(h) HCN
(d) BF
(i) CH
3
3
NH
2
(e) CO
(j) N
2
2
O
_________________________________________________________________________
10.2 The Properties of Liquids
Surface tension is the force that holds molecules on the liquid surface and makes the surface appears tight like a this film. Surface tension is due to cohesive (attractive) forces between molecules at the surface and those under the surface. This net "inward" pulls of molecules on the surface cause the liquid surface to curve. Liquids with strong intermolecular forces have high surface tension and tend to form spherical droplets, which minimize the surface area of the liquid. For example, strong hydrogen bonds in water are responsible for the high surface tension of water, which tends to form spherical droplets. Surface tension is expressed as energy (in Joules) per square meter (J/m 2 ). The surface tension of water is 7.28 x 10 -2 J/m 2 at 20 temperature. o C. Surface tension decreases as temperature increases because of the weakening of the intermolecular forces as molecules move and vibrate faster at high
Liquids with high surface tension, such as water, are not good cleaning solvents. They do not wet the surface to be cleaned effectively. Their molecules cling to each other more strongly rather than with dirt particles on the surface to be cleaned. Adding soap, detergent,
3

or surfactant to water decreases its surface tension and makes the solution a better cleaning solvent. Soap and detergent molecules cause discontinuity in the water matrix by weakening and breaking hydrogen bonds. This allows water to wet surfaces more effectively, and the presence of “micelles” in the soapy solution facilitates the removal of greasy substances from utensils or fabrics. Organic solvents such as acetone, alcohol and kerosene, which have weak cohesive forces and low surface tension, are good cleaning solvents for grease and oily surfaces. Organic solvents also contain nonpolar molecules and they are more suitable for cleaning oily and greasy surfaces.
Capillary action is the tendency of a liquid to rise in narrow and thin capillaries or on absorbing materials. Two types of forces are responsible in capillary action:
- attractions between liquid molecules, and
- attractions between liquid molecules and capillary wall. The relative strengths of adhesive and cohesive forces determine the shape of meniscus of the liquid and how high the capillary will rise. Strong cohesive and adhesive forces cause the high capillary action of water. Strong adhesive forces also cause water meniscus to curve upward. Adhesive forces are dependent on the materials that are in contact with water. For example, adhesive forces between water and woolen materials are much weaker than those between water and cotton fabrics. As a result, woolen materials get wet less readily than cotton fabrics.
Viscosity is the resistance to the flow of liquids. Cohesive forces in liquids create an internal friction that reduces their fluidity. The viscosity of water is 1.0 x 10 -3 kg/m.s at 20 o C, whereas the viscosity of glycerol at this temperature is 1.5 kg/(m.s). Glycerol, honey, and syrup are more viscous than water because they have stronger intermolecular forces, namely, dispersion (London) forces and hydrogen bonds. Like surface tension, viscosity decreases as temperature increases.
Evaporation occurs because molecules on liquid surface acquire sufficient kinetic energy to overcome intermolecular attractions. Evaporation is an endothermic process.
Condensation, the reverse process, is an exothermic process. As some liquid molecules escape into the vapor phase, the average kinetic energy of molecules left behind in the liquid decreases and its temperature drops. This causes heat from the surroundings to flow into the liquid, making the surrounding cooler. Sweating is a cooling mechanism for the body - as sweat evaporates from the body it cools the body by removing some of the body heat through evaporation. The rate of evaporation increases as the surface area and temperature increase.
Exercise-2:
1. Which compound, CH
3
CH
2
OH or HOCH
2
CH
2
OH, has the higher surface tension? Explain.
2. Which compound, CH
3
CH
2
OH or CH
3
CH
2
Cl, has the higher boiling point? Explain.
3. Which compound, CH
3
OH or CH
3
CH
2
OH, has the higher enthalpy of vaporization? Explain.
4. Which compound, CH
3
Cl or CH
3
Br, has the higher vapor pressure at a given temperature? Explain.
5. Which compound, CH
3
CH
2
CH
2
OH or HOCH
2
CH(OH)CH
2
OH, has the higher viscosity?
Explain.
__________________________________________________________________________
4

10.3 Structures and Types of Solids
In the broadest sense, solids are divided into two categories –
and
solids. Crystalline solids have regular arrangements or patterns of lattice structures of the components (atoms, ions, or molecules); amorphous solids have their components arranged in a rather disorderly manner. Common glass is an example of amorphous solid.
Many characteristic shapes of crystalline solids, such as the diamond structure and the rock salt structure are due to the regular arrangement of the particles called
in the crystals.
The smallest repeating unit containing a minimum number of lattice points that may represent the entire crystal is called a
. The entire crystal can be generated by repeating its unit cell in all three dimensions.
There are many types of crystalline solids. They are classified by the types and lattice patterns of components that make up the substance and how these components (atoms, ions, or molecules) are bonded together. The following classifications for crystalline solids may also apply to amorphous solids.
 Molecular solids - solids that contain molecules: ice, dry ice, sugar, etc.
 Ionic solids - solids that contain ions: all salts;
 Atomic solids – metallic, covalent network, and Group 8A solids.
Molecular solids contain molecules that are held together by weak intermolecular forces such as dispersion forces, permanent dipole-dipole attractions, and hydrogen bonding. For examples, in ice, water molecules are held together by hydrogen bonds; in dry ice (solid CO
2
) and iodine molecules are held together by London dispersion forces. These intermolecular forces are relatively very weak forces and molecular solids are characterized by low melting
points. They are also nonconductors.
All ionic compounds form ionic crystals – they contain ions held firmly together by strong ionic bonds. Ionic solids are characterized by high melting points. For example, NaCl melts at about 801 o C. Ionic solids cannot conduct electricity because the ions are not free to flow. However, when melted or dissolved in water, the ions are free to flow and they form excellent conductors.
Atomic solids are divided into three subgroups: metallic solids, covalent network solids, and solids formed by the noble gas (Group 8A) elements. All metals form metallic solids, which are characterized by high thermal and electrical conductivity, malleability, and ductility.
In metallic solids atoms are held together by a “sea” of delocalized valence electrons, which make metals an excellent thermal and electrical conductors. The “sea-of-electrons” model also explains why metals are malleable and ductile. The layers of atoms that form metallic crystals are held firmly by the “sea” of valence electrons, which act like a glue. These layers of atoms can be forced to slide or glide over each other without breaking. The interaction of delocalized electrons with light also accounts for the lustrous appearance of metals.
Metals have variable hardness depending on the number of valence electrons holding the atoms together. For example, Group 1A metals are generally soft because they have the least number of valence electrons. Most transition metals are relatively hard because they have more electrons in the ns and (n-1)d subshells, which make up the "sea of electrons”.
5

Covalent network solids such as diamond and silica, contain atoms that are covalently bonded to each other in a network pattern. They are very hard and have extremely high melting points. Diamond is the hardest known natural substance with melting point of approximately 3600 o C. They are nonconductors. Graphite is a semi-covalent network solid, which is made of layers of carbon atoms; within each layer carbon atoms are covalently bonded, but the forces holding these layers of carbon atoms are weak van der Waals forces.
These weak forces allow the carbon layers to slide over each other and can flake off. Thus graphite is much softer than diamond although both are composed of carbon atoms. Graphite is used as pencil lead and lubricant, whereas diamond is used as glass cutting utensil and in rock drill bits. Unlike diamond, graphite can conduct electricity because it contains delocalized
 -electron system.
In the Group 8A solids, noble gas atoms are held together by weak dispersion forces.
As a result, this type of atomic solids has very low melting points. Unlike metallic solids, this type of solids is a nonconductor because their valence electrons are not delocalized.
10.4 Structure and Bonding in Metals
Metallic crystals can be pictured as containing spherical atoms packed together in regular lattice and are held together by a “sea” of delocalized valence electrons. Metals crystallize either into hexagonal closest packing, cubic closest-packing (or face-centered cubic), body-centered cubic, or simple cubic structures. The simple cubic structure is the least common in metals.
and
(FCC)
In the closest-packing arrangement, each sphere has 12 nearest neighbors (6 on the same layer, 3 above, and 3 below that layer). Layers of spherical atoms are stacked in such a way that each atom of one layer is nested in the depression of the two adjacent layers.
When spheres in the second layer of a close-packed arrangement are nested on top of the first layer, two types of “holes” are formed in the second layer. These are the tetrahedral holes - those directly on top of spheres in the first layer, and octahedral holes - those directly on top of “holes” in the first layer. Sphere of the third layer may be nested on top of the
tetrahedral holes, which will result in an ABABAB…pattern of arrangement of spheres, or on top of the octahedral holes, which yields an ABCABC.. pattern.
The closest-packed arrangement with the ABCABC.. pattern forms the face-centered
cubic (fcc) or cubic close-packed structure. While the ABABAB.. pattern is called the
hexagonal closest-packed (hcp) structure. In both types of arrangements, the maximum space occupied by spheres is about 74% (26% is void space).
(BCC)
In the body-centered cubic structure, each sphere has the coordination number 8 (or eight nearest neighbors). The body-centered cubic structure has a higher percentage of occupied space than the simple cubic structure. The maximum occupied space in a bodycentered cubic arrangement is about 69%. Many metals crystallize into the body-centered cubic structure, such as chromium, manganese, and iron.
(SC)
In the simple cubic structure spheres (atoms or ions) are arranged in layers, one placed directly on top, below, or next to the other. Each sphere has 6 nearest neighbors, which is also called the
. In the simple cubic structure, the maximum space occupied by spheres is only 52.36% of the available space (47.64% is void space).
6

The structures of crystalline solids are determined by the manner atoms are arranged and the resulting lattice patterns of the unit cells. All cubic unit cells have a lattice point at each corner of the cube, which are the only lattice points in the simple cubic unit cell. In the
body-centered cubic (bcc) unit cell an additional lattice point occurs at the center of the cube; while the face-centered cubic (fcc) unit cell has a lattice point at the center of each face of the cube in addition to the eight corner positions.
Spheres at corner and face-center positions are shared by neighboring unit cells, but one in the center of the cube is not. Thus, for a given unit cell in a crystal, only one-eighth of a sphere at each corner and one-half of a sphere on each face-center belongs to that unit cell.
The whole sphere at the center of the cube belongs to that unit cell. The total number of
"complete" spheres in each type of cubic unit cell is as follows:
 Simple cubic: total number of spheres = (8 x  ) = 1;
 Body-centered cubic: total number of spheres = (8 x  ) + 1 = 2;
 Face-centered cubic: total number of spheres = (8 x  ) + (6 x ½) = 4.
Exercise-3
1. Platinum crystallizes into a cubic lattice structure that has a total of four Pt atoms per unit cell. (a) What type of unit cell does platinum form? (b) What is the mass of a unit cell of platinum? (c) If the density of platinum is 21.45 g/cm of the unit cell and the atomic radius of platinum, respectively. Express the edge length and atomic radius in picometers. (1 u = 1.6605 x 10 -24
3 , what are the edge length g; 1 pm = 10 -10 cm)
2. Iron forms a crystal lattice with body-centered cubic unit cell. If the atomic radius of iron is 126 pm, what is the measurement of the side length of the unit cell of iron? If the mass of iron atom is 55.847 u, calculate the density of iron. (1 u = 1.6605 x 10 -24 g)
3. Gold forms a face-centered cubic unit cell lattice and has a density of 19.32 g/cm
Calculate the radius of a gold atom.
3 .
_________________________________________________________________________
Bonding model for metals must account for the physical properties of metal, especially their efficient conductivity, malleability, and ductility. The simplest picture that explains these properties is the “electron sea model”. It depicts metals as containing a regular array of cations in a “sea” of delocalized valence electrons. The mobile electrons can conduct heat and electricity, and the metal ions can be moved around as the metal is hammered in to a sheet or pulled into a wire without breaking it.
The molecular orbital theory proposes a band model to explain the conductivity of metals. Note that when two atoms combined to form a diatomic molecule, their valence atomic orbitals combine to form two set of molecular orbitals – the bonding molecular orbitals and antibonding molecular orbitals, with widely spaced energy levels. If there are n atoms combine in these manner, there will be 2n sets of these molecular orbitals, but with their energy levels become closely spaced and form a virtual continuum of levels, called bands.
Valence electrons are assumed to travel around the metal crystal in these partially filled bands.
7

For example, in magnesium metal, which forms hcp structure, the 3s and 3p orbitals of the valence shells from all the atoms that make up the crystal combine to form the 3s and 3p bands. The 3s-band in magnesium is filled band, but the 3p-band is empty. Because the two valence bands overlap one another, electrons from the filled 3s-band can move into the 3pband and move around freely over the entire crystal structure. The 3s and 3p bands in magnesium form the conduction band. In transition metals the conduction band would include the (n – 1)d, ns, and np bands.
Alloys are homogeneous mixture of two or more metals. There are two types of alloys
–
and
. In substitutional alloys, some of the host atoms are replaced by atoms of another metal. This type of alloys normally involves metals that have atoms of comparable size. For example, brass is a substitutional alloy in which approximately a third of the (host) copper atoms are replaced with zinc atoms. Other substitutional alloys are
Sterling silver (93% Ag, 7% Cu) and pewter (85% Sn, 7% Cu, 6% Bi, and 2% Sb).
An interstitial alloy is formed when atoms of another element are small enough to occupy interstitial holes formed by the larger atoms in the host crystal. Carbon steel is an example of interstitial alloy, in which the small carbon atoms occupy some of the interstitial holes in the iron crystal. The presence of carbon atoms in the interstitial holes of the iron crystal makes the resulting steel harder, stronger, and less malleable than pure iron. The type, strength, and hardness of the steel depend on the amount of carbon in the steel.
However, too much carbon will make the steal brittle. Mild steels contain less than 0.2% carbon (by mass) and they are quite malleable and ductile. They are used to make nails, cables and chain-links. A medium steel contains 0.2 – 0.6% carbon and is harder than mild steels. They are used to make structural steel beams and railway tracks. High-carbon steels, which contains 0.6 – 1.5% carbon, are very hard and are used to make springs and strong cutlery.
Some familiar alloys, their compositions and Uses.
——————————————————————————————————————
Name Compositions (mass %) Uses
——————————————————————————————————————
Stainless steel 73-79% Fe, 1q4-18% Cr, 7-9% Ni Cutlery, instruments
Nickel steel
High speed steel
96-98% Fe, 2-4% Ni
80-84% Fe, 14-20% W
Cables, gears
Cutting tools
Perm alloys
Bronze
Brass
78% Ni, 22% Fe
70-95% Cu, 1-25% Zn, 1-18% Sn
50-80% Cu, 20-50% Zn ocean cables
Statues, castings
Plating, ornamental objects
Sterling silver
14-carat gold
92.5% Ag, 7.5% Cu
58% Au, 4-28% Ag, 14-20% Cu
Jewelry
Jewelry
Dental amalgam 69% Ag, 18% Sn, 12% Cu, 1% Zn dental fillings
——————————————————————————————————————
10.5 Covalent Network Atomic Solids – Carbon and Silicon
Elements like carbon (in diamond) and silicon form crystals in which atoms are covalently bonded to each other. This type of solids is very hard and is nonconductors. In diamond each carbon atom is covalently bonded to four other carbon atoms that involves sp hybridized atomic orbitals on each carbon atom. Since all valence electrons are involved in
3 covalent bonding, there is no delocalized valence electron in diamond, which makes diamond a non-conductor. The energy gap between filled bonding MOs and empty antibonding MOs are
8

too large to allow electrons from the lower filled MOs to jump the energy gap to the higher energy unoccupied MOs and become delocalized.
In the graphite structure, each carbon atom has sp 2 hybridized atomic orbitals, which are involved in covalent bond formation with three other carbon atoms. Since the geometry about each carbon is trigonal planar, the entire networks are composed of flat sheets of carbon atoms, within which atoms are covalently bonded. However, interactions between these carbon layers are weak van der Waals forces, which allow graphite layers to slide over each other. Thus graphite is much softer than diamond. Additionally, in sp 2 hybridization, each carbon atom also has an unhybridized 2p-orbital that contains an electron. These unhybridized orbitals combine to form delocalized  -electron system over the entire graphite structure. This enables graphite to conduct electric current.
Because of their extreme hardness, diamonds are used extensively in industrial cutting implements. Since natural diamonds are very expensive and rare, many industries turn into synthetic diamonds, which are formed by subjecting graphite to a pressure above 1.5 x 10 atm at temperature exceeding 2800 o
5
C. A thin layer of such synthetic diamonds on the glass surface makes the latter scratch-resistance.
Silicon is the second most abundant element in the Earth crust, found mainly in sand and rocks. Although silicon belongs to the same group as carbon, their compounds are quite different. For example, carbon and oxygen form discrete molecules like CO and CO
2
, but silicon does not form such molecules. Silica (sand) and quartz have the empirical formula
SiO
2
, but they are covalent network solids with structures based on a network of SiO
4 tetrahedra with shared oxygen atoms rather than discrete SiO
2
molecules. Silicon atoms are too large to form effective sidewise 2p-3p orbital overlaps and forms  -bonds with oxygen atoms.
When sand is heated to above its melting point (about 1600 o C) and then cooled rapidly, an amorphous solid called
is formed. Common glass is formed by adding sodium carbonate, Na
2
CO
3
, to the molten sand and the mixture immediately cooled. The properties of glass can be improved by adding certain additives, such as Al
2
O
3
, B
2
O
3
, Na
2
O, K
2
O, MgO, and
CaO. Some common types of glass and their compositions are given in the following table.
——————————————————————————————————————
Percentages of Components_______________
Types of Glass SiO
2
B
2
O
3
Na
2
O MgO Al
2
O
3
K
2
O CaO
——————————————————————————————————————
Window (soda-lime) glass 72 --- 13 --- 0.3 3.7 11
Cookware (aluminosilicate glass) 55 --- --- 10 20 --- 15
Heat-resistant borosilicate glass 76 13 5 --- 2 0.5 3
Optical glass 69 0.3 6 --- --- 12 12
——————————————————————————————————————
The major compounds of silicon found in rocks, soils, and clays are the
, which structures are based on interconnected SiO
4
units. Unlike silica, silicates are ionic compounds containing metal cations and polyatomic anions composed of silicon and oxygen. The properties of silicates depend on the size of the “SiO clusters. For example, calcium orthosilicate, Ca
2
SiO
4
4
” units forming the polyatomic anion
, is a component of the Portland cement.
Ceramics
Ceramics are made from clays and hardened by firing at high temperatures. They are nonmetallic materials that are strong, brittle, and resistant to heat and chemicals. Like glass,
9

ceramics are silicate-based materials, but unlike glass ceramics cannot be re-melted once they are hardened by firing at high temperature. Whereas glass is a homogeneous, noncrystalline substance, a ceramic is a heterogeneous solid, which contains two phases: fine crystals of silicates suspended in a glassy cement. The formation of ceramics is based on the properties of clays, which are formed by the weathering action of water and carbon dioxide on the mineral feldspar. The latter is a mixture of silicates with empirical formulas such as
K
2
O .
Al
2
O
3
.
6SiO
2
and Na
2
O .
Al
2
O
3
.
6SiO
2
. The weathering of feldspar produces kaolinite, which is an aluminosilicate consisting of tiny thin platelets with the empirical formula Al
2
Si
2
O
5
(OH)
4
.
When dry, the platelets form interlocking structures. When the remaining water is driven off by firing, the silicates and cations form a glass that binds the tiny crystals of kaolinite.
Semiconductors
Elemental silicon has the same structure as diamond. However, unlike diamond that has a large energy gap between filled and empty MOs, the energy gaps between filled MOs and empty MOs in silicon are much smaller. Even at 25 o C some electrons can be excited into the empty MOs and become delocalized. As temperature increases, more electrons can be excited and become delocalized, which increases the conductivity of the materials. Thus, the conductivity of silicon (a semi-metal) increases with temperature. In contrast, the conductivity of metals decreases as their temperature increases.
The conductivity of silicon can also be enhanced by doping the silicon crystal with elements such as boron or arsenic. Doping silicon with arsenic produces an
, in which the extra valence electrons from arsenic that are not involved in covalent bond formation can flow through the crystal when an electric field is applied. Doping silicon with boron produces empty bonding orbitals referred to as “positive holes” into which neighboring electrons can flow. When electric field is applied, these “positive holes” will appear to flow in the crystal. Thus, the doping produces a
.
The most important application involving n- and p- type semiconductors is the formation of a
. The n-type semiconductor contains an “excess” of electrons not involved in bonding, while the p-type semiconductor contains an “excess” of empty orbitals or
“positive holes”. At the junction, a small number of electrons from the n-type flow into the
“positive holes” in the p-type. Such electron migration causes the p-type region to become negatively charged because it contains a surplus of electrons. While the n-type becomes positively charged since it has lost some of its electrons. Such charge build-up at the junction is called the contact potential or junction p[potential, and it prevents further flow of electrons from the n-type to the p-type semiconductors.
If the negative terminal of an electrical source is connected to the p-type region and the positive terminal to the n-type region, electric current will flow in the direction opposite to the natural flow of electrons (and holes) at the p-n junction. The junction will resist the imposed current in this direction, which is said to be under reversed bias. No current will flow in the circuit.
On the other hand, if the negative terminal of the power source is connected to the ntype region and the positive terminal to the p-type region, the flow of current in the circuit will be in the same direction as the movement of electrons and “positive holes” at the junction.
The junction is said to be under forward bias. A p-n junctions are found in rectifier, a device that converts an alternating current to a direct current.
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10.6 Molecular Solids
Many types of solids contain discrete molecular units occupying the lattice positions in the crystals. For example, in ice and snowflakes, the lattice positions are occupied by water molecules. Other examples of molecular solids include dry ice, which contains CO
2
molecules, yellow sulfur, which contains S
8 and iodine, which contains I
2
molecules, white phosphorus, which contains P
4
molecules,
. These substances are characterized by weak intermolecular forces holding those molecules in their lattice position. As a result the solids have relatively low melting points
10.7 Ionic Solids
The crystal structures of many ionic compounds are built from either simple cubic or face-centered cubic lattices of the anions, while the cations occupy the holes created by the anions. In the simple cubic unit cell, the only available hole is the one at the cell center. Ionic compound such as cesium chloride, CsCl, adopts this lattice structure - simple cubic lattice structure for Cl crystal, the lattice consists of Cl
Na +
ions with Cs + ion occupying the center of the unit cell. In sodium chloride
ions in a face-centered cubic arrangement, while the smaller
ions fill all the octahedral holes created by Cl ions. Note that, in ionic crystals, the description of lattices usually refers to the arrangement of anions, which are the larger ions in the crystal lattice.
In summary, ionic compounds with the general formula MX often occur in either of two common crystal structures: (1) M n+ ions occupy all the cubic holes of a simple cubic X lattice, such as in CsCl and CsBr, or (2) M n+ n-
ions occupy all the octahedral holes of a facecentered cubic X n lattice, as in NaCl. Ionic compounds with the face-centered cubic arrangement, or called the “rock-salt” structure are: all alkali metal halides (except CsCl,
CsBr, and CsI), all oxides and sulfides of the alkaline earth metals, and all oxides of transition metals of the fourth period with the general formula MO.
Exercise-4:
1. The ionic radii of the Na + and Cl ions are 95 pm and 181 pm, respectively. Sodium chloride forms a face-centered cubic lattice of Cl ions, with the Na + ions occupying the octahedral holes. What is side measurement of a unit cell of NaCl in pm? Calculate the density of NaCl in g/cm 3 .
(Atomic masses: Na = 22.99 u; Cl = 35.45 u; 1 u = 1.66 x 10 -24 g; 1 pm = 10 -10 cm)
2. The ionic radii of Cs + and Br ions are 181 pm and 195 pm, respectively. Cesium bromide forms a simple cubic crystal lattice of Br
Calculate the density of CsBr in g/cm 3
with a Cs + ion at the center of cube.
. (Atomic masses: Cs = 132.9 u, Br = 79.90 u)
__________________________________________________________________________
10.8 Vapor Pressure and Change of State
When a liquid is placed in a closed container, both evaporation and condensation occur simultaneously and continuously. When the space above the liquid is saturated with its vapor, the rates of evaporation and condensation are equal and a state of equilibrium is established.
The pressure exerted by the saturated vapor is called
of the liquid, which is influenced by temperature and the strength of intermolecular forces in the liquid.
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 Increasing the temperature will increase its vapor pressure
 Strong intermolecular forces tend to reduce its vapor pressure – liquids with strong intermolecular forces generally have low vapor pressure.
Liquids with weak intermolecular forces evaporate more readily - more molecules escape into the vapor phase under a given condition, which results in high vapor pressure. Liquids with strong intermolecular forces do not evaporate very readily - fewer molecules escape into vapor, resulting in low vapor pressure. That is, strong intermolecular forces result in low vapor pressure and weak intermolecular forces result in high vapor pressure.
For example, at 25 o C water has a vapor pressure of 24 torr, whereas acetone (C
3
H
6
O) and diethyl ether (C
4
H
10
O) have vapor pressures of about 230 torr and 530 torr, respectively.
Since water has stronger hydrogen bonds, its vapor pressure is relatively low. Whereas in acetone and diethyl ether, which have weak dipole-dipole attractions, evaporation occurs much more readily, resulting in higher vapor pressures.
The rate of evaporation and vapor pressure increase with temperature. When the vapor pressure is equal to the external (atmospheric) pressure, the liquid will boil. Thus, the
of a liquid is the temperature at which its vapor pressure is equal to outside pressure.
At the normal boiling point, the vapor pressure of a liquid is 760 torr. The vapor pressure of water at 100 o C is 760 torr; if the external pressure is 1 atm, water boils at 100 o C.
If the pressure exerted on the water surface is less than 1 atm, it will boil at a lower temperature. For example, in Denver, Colorado, where the altitude is about 1600 m and the atmospheric pressure is about 630 torr, water boils at about 95 o C. At the top of Mt. Everest, where the atmospheric pressure is only about 230 torr, the boiling point of water barely reaches 70 o C. In a pressure cooker, the boiling point of water can be raised to about 120 o C by applying an additional 15 psi above atmospheric pressure. Water has low vapor pressure and high boiling point because of the extensive hydrogen bonding.
Liquids like acetone and diethyl ether, which have weak dipole-dipole attractions, have high vapor pressures at 25 o C and low boiling points. Acetone and diethyl ether boil at about
56 o C and 35 o C, respectively. They are called
liquids.
The linear relationship between vapor pressure and temperature is given by the
equation:
P = - (  H vap
) . 1 + C;
R T
(P
2
) = - (  H vap
) ( 1 - 1 ) = (  H vap
) (T
2
- T
1
)
P
1 where,  H vap
R T
2
T
1
R T
1
T
2
is the molar enthalpy of vaporization for the liquid, which is assumed to be constant over the range of temperature T
1
to T
2
; R = 8.314 J/(mol.K).
,  H vap
, is the quantity of heat needed to evaporate a
mole of liquid at 1 atm and constant temperature.
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Exercise-5:
1. The boiling point of iso-octane, C
8
H
18
, is 99.2
o C and its enthalpy of vaporization,  H vap
=
35.76 kJ/mol. Determine the vapor pressure of iso-octane at 25 o C.
2. If hexane, C
6
H
14
, has an enthaply of vaporization,  H vap
= 30.1 kJ/mol, and its vapor pressure at 25 o C is 148 mmHg, determine the boiling point of hexane.
3. The melting point of potassium is 63.2
o C. Molten potassium has a vapor pressure of 10.0 torr at 443 o C and a vapor pressure of 400.0 torr at 708 o C. (a) Use these data and
Clausius-Clapeyron equation to calculate the heat of vaporization of liquid potassium.
(b) Calculate the vapor pressure of liquid potassium at 100 o C.
___________________________________________________________________________
The melting point (or freezing point) is the temperature at which solid and liquid are at equilibrium under a given pressure. The
(or
) of a liquid is
. For example, ice and liquid water are at equilibrium and have the same vapor pressure at 0 o C and 1 atm. Under these conditions, the rates of ice melting and water freezing are equal. Melting points vary very little with changes in external pressure.
A heating curve is a graph of relationship between temperature and heat absorbed.
Two constant temperature regions are observed in such curves - one at the melting point of solid and another at the boiling point of liquid. When a substance occurs as a single state
(solid, liquid or vapor), its temperature increases gradually as heat is absorbed. Temperature remains constant as long as both solid and liquid occur together. As soon as the solid has melted completely, the temperature of the resulting liquid will rise again gradually as more heat is absorbed until the boiling point is reached. At the boiling point, both liquid and vapor are at equilibrium. Once again the temperature remains constant until the liquid has vaporized completely.
During melting and evaporation, heat is absorbed without raising the temperature of the substance. Heat absorbed during melting is called the enthalpy of fusion,  H f
. Heat absorbed during evaporation is called the enthalpy of vaporization,  H vap
, which is the amount of heat needed to break all intermolecular forces that occur in the liquid state. The enthalpy of vaporization of a substance is normally about 5 - 6 times greater that its enthalpy of fusion.
A cooling curve is the reverse of the heating curve. It shows the decrease in temperature as a function of heat lost. The two constant temperature regions occur at the same temperature values as in the heating curve. If the cooling process is carried out very rapidly, a liquid can be super cooled to a temperature well below its normal freezing point, but does not solidify. This is called a super cooled liquid, which is an unstable state. “seeding” with a tiny crystal or even a speck of dust can easily bring about crystallization of a super cooled liquid.
,  H fus
, is the amount of heat required to melt one mole of a
substance at its melting point.
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Exercise-6:
1. (a) 933 J of heat is required to vaporize 1.75 g of acetone, C
3
H
6
O, at 298 K. What is the enthalpy of vaporization,  H vap
(in kJ/mol), of acetone at 298 K? (b) Using the  H vap calculated in part (a) and the normal boiling point of 56.5
o vapor pressure at 25 o C.
C for acetone, estimate its
2. 23.0 g of ice is melted at 0 vaporized at 100
 H o fus
= 6.02 kJ/mol;  H o C. The water is then heated to 100 o C and then completely
C. Calculate the total amount of heat (in kJ) absorbed by the system. vap
= 40.7 kJ/mol; Sp. heat = 4.18 J/g.
o C.
3. (a) 5.0 g of water is place in a 5.0-liter flask from which air has been completely evacuated. How much of the water will vaporize at 25
25 o o C? The vapor pressure of water at
C is 23.8 torr. (b) How much heat is absorbed in evaporation? (  H vap
= 44.0 kJ/mol)
___________________________________________________________________________
A phase diagram is a graph that summarizes the pressure-temperature relationships associated with
and the changes of the phases of a substances. The phase diagram for a single substance, such as water, contains three areas - one for each phase
(solid, liquid and vapor). Lines or curves in the graph represent pressure-temperature relationships where two phases (solid and liquid, liquid and vapor, or solid and vapor) are at equilibrium. The three phases meet at the
- a point at which all three phases are at equilibrium. The triple point for water occurs at 0.0098
water and water vapor are at equilibrium. o C and 4.58 torr, at which ice, liquid
Water has a unique phase diagram. Unlike other substances, the solid/liquid boundary line for water has a negative slope instead of a positive one like most substances. This means that the melting point of ice decreases if an external pressure is applied on it. This behavior is associated with another property of water, which is that its density increases as ice melts. In fact, water reaches a maximum density of 1.00 g/mL at about 4 o C. Most other substances decrease in density as the solids melt. Since ice is less dense than water, it floats and this has special implication. For example, during cold Winter season, some lakes become frozen. The top layer of the lake water freezes first because the direct contact with the cold air. The ice formed on the top layer insulates the bottom layer of the lake and prevent further freezing.
This allows aquatic life, especially fish to survive through the Winter season. If ice has been denser than water, it would sink and the entire lake will be frozen and no aquatic life will survive the Winter season. When water freezes, it volume expands, which explains for the decrease in the density during freezing. In countries or regions where Winter is extremely cold, water pipes need to be insulated to prevent the water in the pipe from freezing, which can cause the pipe to crack or burst as a result of volume expansion.
In the phase diagram of carbon dioxide, the solid-liquid boundary line has a positive slope, which is typical of most substances. At the solid-liquid equilibrium, the formation of solid CO
2
is favored at high pressure, which is the opposite to that of water. Solid CO
2
is also denser that the liquid form. The triple point for carbon dioxide occurs at 5.1 atm and –56.6
and its critical point is at 72.3 atm and 31 o C. This implies that liquid CO
2
will only form at o C, pressure above 5.1 atm and temperature below 31 o temperature above –78 o C, solid CO
2
C. At 1 atmospheric pressure and
becomes gas directly. Thus the name “dry ice”. In CO fire extinguishers, the gas is under high pressure and carbon dioxide is a liquid. When the
2
14

pressure is released, it immediately vaporizes. Since CO
2
gas is heavier than air, it forms a blanket over the fire and smothers it. Since vaporization is an endothermic process, it also has a cooling effect.
If a liquid is heated in a sealed capsule, its vapor (and vapor pressure) will build up on the surface above the liquid. A point will be reached where its meniscus just disappears and it appears neither as real liquid nor vapor. This point is called a
, which is a temperature and pressure at which liquid-vapor boundary disappears. The critical point for water occurs at 374 o C and 218 atm. At or above the critical point, a substance exists as
or plasma state. Carbon dioxide form supercritical fluid at 72.3 atm and
31 o C. Above the critical temperature, a gaseous substance cannot be liquefied regardless of the pressure applied upon it. For example, the critical point for nitrogen occurs at 126 K (-
147 o C) and 33.5 atm. Nitrogen gas can only be liquefied at temperature below 126 K. Above this temperature, it will never becomes liquid, regardless of the applied pressure.
15