2012 General Chemistry I 1 Chapter 5. LIQUIDS AND SOLIDS 2012 General Chemistry I INTERMOLECULAR FORCES LIQUID STRUCTURE 5.1 The Origin of Intermolecular

2012 General Chemistry I 1

Chapter 5. LIQUIDS AND SOLIDS

2012 General Chemistry I

INTERMOLECULAR FORCES

LIQUID STRUCTURE

5.1 The Origin of Intermolecular Forces5.2 Ion-Dipole Forces5.3 Dipole-Dipole Forces5.4 London Forces5.5 Hydrogen Bonding5.6 Repulsions

5.7 Order in Liquids5.8 Viscosity and Surface Tension


INTERMOLECULAR FORCES (Sections 5.1-5.6)

5.1 The Origin of Intermolecular Forces5.1 The Origin of Intermolecular Forces

Phase: uniform in both chemical composition and

physical state

- Condensed phase: simply a solid or liquid phase

- Condensed phases form when attractive intermolecular forces between molecules pull them together; repulsions dominate at even shorter separations.


- Intermolecular forces are weak compared with bonding forces, but boiling points and sublimation points depend on their strength.

- All interionic and almost all intermolecular forces can be traced to the coulombic interaction between charges.


1/r : between ions (ionic bonding)

- Distance dependence of potential energy of interaction

1/r2 : between ions and dipoles1/r3 : between stationary dipoles1/r6 : between rotating dipoles


TABLE 5.1 Interionic and Intermolecular Interactions


5.2 Ion-Dipole Forces5.2 Ion-Dipole Forces

Hydration: attachment of water molecules to ionic solute particles is an example of ion-dipole interaction.

H2O

This is the attractive force between ions and polar molecules in liquid or solid phase.


The potential energy of ion-dipole interactions (~15 kJmol-1)

z = the charge number of the ion = the electric dipole moment of the polar molecule


Water of crystallization: smaller and highly charged cations strongly attract polar water

molecules in the solid phase.

- Note hydrated salts of Li and Na vs. anhydrous salts of K, Rb, Cs, and NH4

+

- Note BaCl2 · 2H2O vs. anhydrous KCl (Ba2+; 136 pm, K+; 138 pm)

E.g. Na2CO3.10H2O compared with K2CO3

(Na+; 102 pm, K+; 138 pm)


5.3 Dipole-Dipole Forces5.3 Dipole-Dipole Forces

Between stationary polar molecules in the liquid phase (~2 kJ mol-1)

Between rotating polar molecules in the gas phase (~0.3 kJ mol-1)

This is the attractive force between polar molecules.


Self-Test 5.1B

Which will have the higher boiling point, 1,1-dichloroetheneor trans-1,2-dichloroethene?

Solution

1,1-dichloroethene is polar, whereas trans-1,2-dichloroetheneis nonpolar:

C CCl

Cl

H

HC C

Cl

H

H

Cl

Hence, dipole-dipole (as well as London) forces existin 1,1-dichloroethene, giving it the higher boiling point.


5.4 London Forces5.4 London Forces

London force (induce dipole-induced dipole force, dispersion force) ~2 kJmol-1: it exists between all molecules but is the only interaction between nonpolar molecules.

- Attractive interactions due to instantaneousfleeting dipole moments

- Fluctuation of the electron distribution in one molecule→ temporary dipole → secondtemporary dipole in the other molecule → ···

Time


Polarizability is the ease with which molecular electron clouds can be distorted: increases with number of electrons.

- Potential energy (strength) of the London interaction is given by

- A large linear molecule is more likely to have stronger London interactions (and hence a higher boiling point) than a smaller or nonlinear one.

C5H12; mobile liquid C15H32; viscous liquidC18H38; waxy solid

Examples

Alkanes


- Halogens: gases (F2, and Cl2); liquid (Br2); solid (I2)

- Rod-shaped (pentane; Tb = 36 oC) vs. spherical (2,2-dimethylpropane; Tb = 10 oC)


TABLE 5.2 Melting and Boiling Points of Substances


Dipole-induced dipole interaction between a polar molecule and a nonpolar molecule (~2 kJ mol-1)

- Dipole-dipole force between rotating polar molecules, London force, dipole-induced dipole force

Van der Waals interactions

Allied intermolecular interactions


EXAMPLE 5.2

Explain the trend in the boiling points of the hydrogen halides:HCl, -85 oC; HBr, -67 oC; HI, -35 oC.

- Electronegativity differences: HCl > HBr > HI

- Number of electrons and London forces: HCl < HBr < HI

→ not by dipole-dipole forces, but by London forces

Self-Test 5.2A

Account for the trend in boiling points of the noble gases,which increase from helium to xenon.

Solution

In the noble gases, only London forces need be considered:These increase as the number of electrons increases (sizeof the atom increases):

He (2)

B.Pt

Ne(10) Ar(18) Kr(36) Xe(54)


5.5 Hydrogen Bonding5.5 Hydrogen Bonding

Some compounds are characterized by exceptionally high Tb due to hydrogen bonding: examples include NH3, H2O, HF – see Fig. 5.9.

London forces

Hydrogen bonding


Hydrogen bonding: an attraction in which a hydrogen atom bonded to a small, strongly electronegative atom, specifically N, O, or F, is attracted to a lone pair of electrons on another N, O, or F atom. Intermolecular and intramolecule types exist.

- strong electrostatic interaction ~20 kJmol-1

O…….H-O

linear but asymmetric (101 pm vs 175 pm)


- Hydrogen fluoride, (HF)n - Acetic acid dimer (vapor)

- Protein folding

- DNA doublehelix


Self-Test 5.3B

Which of the following molecules can take part in hydrogenbonding with other molecules of the same compound:(a) CH3OH; (b) PH3; (c) H-O-Cl?

Solution

(a) and (c), since these both have a H-O covalent bond:

O

H3C

H

H O

CH3

O

Cl

H

H O

Cl

(most likely)


5.6 Repulsions5.6 Repulsions

-Intermolecular repulsions arise from the overlap of orbitals on neighboring molecules and the requirements of the Pauli exclusion principle.- They are important only at very short distances:

Ep

1

r12


LIQUID STRUCTURE (Sections 5.7-5.8)

5.7 Order in Liquids5.7 Order in Liquids

- The liquid phase lies between the extremes of the gas and solid phases.

gas phase: moving with almost complete freedomminimal intermolecular forces

solid phase: locked in place by intermolecular forcesoscillate around an average location


- In the liquid phase, molecules have short-range order but not long-range order.

- Water loses only 10% of hydrogen bonds uponmelting and the rest are continuously brokenand reformed.


5.8 Viscosity and Surface Tension5.8 Viscosity and Surface Tension

Viscosity: resistance to flow, indication of the intermolecular force strength

- Water and glycerol: very viscous due to hydrogen bonding

- Hydrocarbon oils and grease: viscous due to tangling long chains

- Viscosity usually decreases with temperature due to higher energy of molecules.


Viscosities of common liquids

Linear alkane chains inHeavy hydrocarbon oil


Surface tension: the net inward pull, an indication of the intermolecular force strength

- water: three times larger than many other

liquids, due to hydrogen bonds

- mercury: more than six times that of water,

partially covalent


Wetting: strong interactions of water with the materials’ surface. Water maximizes its contact with the materials by hydrogen bonding.

Capillary action: adhesive forces between a liquid and surface vs. cohesive forces withinthe liquid

- Meniscus: indication of the relative strength of adhesion and cohesion

H2O Hg


Chapter 5. LIQUIDS AND SOLIDS

2012 General Chemistry I

SOLID STRUCTURES

THE IMPACT ON MATERIALS

5.9 Classification of Solids5.10 Molecular Solids5.11 Network Solids5.12 Metallic Solids5.13 Unit Cells5.14 Ionic Structures

5.15 Liquid Crystals5.16 Ionic Liquids


SOLID STRUCTURES (Sections 5.9-5.14)

5.9 Classification of Solids5.9 Classification of Solids

Crystalline solid: a solid in which the atoms, ions, or molecules lie in an orderly array with crystal faces

Amorphous solid: one in which the atoms, ions, or molecules lie in a random jumble

quartz amorphous silica


Classification of Crystalline Solids According to the bonds that hold their atoms, ions, or molecules in place:

Molecular: assemblies of discrete molecules held in place by intermolecular forces

Metallic: consisting of cations held together by a sea of electrons

Ionic: built from the mutual attractions of cations and anions

Network: consisting of atoms covalently bonded to their neighbors throughout the extent of the solid



5.10 Molecular Solids5.10 Molecular Solids

Molecular solids consist of molecules held together by intermolecular forces; physical properties depend on the strengths of those forces.

Amorphous molecular solids: as soft as paraffin wax

Crystalline molecular solids:

- sucrose: numerous hydrogen bonds between OH

Groups account for high melting point at 184 oC

- ultrahigh-density polyethylene: smooth yet tough


- ice

water benzene

Molecular Solids and Liquids: Melting andFreezing

Most substances increase in density on freezing:water is an important exception. Ice at 0 oC isless dense than water at 0 oC due to a moreopen hydrogen-bonded structure.


5.11 Network Solids5.11 Network Solids

Network solids are characterized by a strong covalent bond network throughout the crystal: they are very hard and rigid, with high Tm and Tb

E.g. Two common allotropes (forms of an element that differ in the way in which the atoms are linked) of carbon have very different network structures.

Diamond, with an sp3 hybrid -bonding framework, is one of the hardest substances


- Ceramic materials: noncrystalline inorganic oxides, great strength

Graphite has flat sheets of sp2 hybrid s-bonds with weak bonding between sheets. It conducts electric parallel to the sheet, and is soft and slippery


5.12 Metallic Solids5.12 Metallic Solids

In metallic solids, the cations are bound together by their interaction with the sea of the electrons that they have lost.

Close-packed structure: the spheres stack together with the least waste of space

- Hexagonal close-packed structure (hcp): packed with the sequence of ABABAB···

Coordination number: the number of nearest neighbors of each atom

Coordination number = 12(3 plane below + 6 own plane + 3 plane above): this is the maximum.


- Cubic close-packed structure (ccp): packed with the sequence of ABCABC··· Coordination number of ccp = 12

Occupied space in a ccp:


Holes: the gaps (interstices) between the atoms in a crystal

- Tetrahedral hole: a dip between three atoms is directly covered by another atom

- Octahedral hole: a dip in a layer coincides with a dip in the next layer


5.13 Unit Cells5.13 Unit Cells

Unit cell: the smallest unit that, when stacked together repeatedly without any gaps and without rotations, can reproduce the entire crystal.

- Face centered cubic (fcc, cubic F)

- Body centered cubic (bcc, cubic I)

- Primitive cubic (cubic P)

Cubic F Cubic I Cubic P


Bravais lattices: 14 basic patterns of unit cell in 3D crystalline Systems; P = primitive; I = body-centered; F = face-centered; C = with lattice point on two opposite faces; R = rhombohedral


Unit cells are characterized bylengths a, b, c and angles , ,

- Cubic unit cells

- Face centered cubic (fcc, cubic F)

- Body centered cubic (bcc, cubic I)

- Primitive cubic (cubic P)




Self-Test 5.4A

How many atoms are there in a primitive cubic cell?

Solution

In a cubic P cell there are only eight corner atoms,hence the total number of atoms per unit cell is:8 x 1/8 = 1.


EXAMPLE 5.3

The density of copper is 8.93 gcm-3 and its atomic radius is 128 pm.Is the metal (a) close-packed or (b) body-centered cubic?

(a) Fcc (ccp) and hcp cannot be distinguished by density only.

For 4 atoms in a fcc cell,


(b) For 2 atoms in a bcc shell,

close-packed cubic (fcc)


Self-Test 5.5A

The atomic radius of silver is 144 pm and its density is10.5 g cm-3. Is the stucture face-centered cubic (fcc;close packed) or body-centered cubic (bcc)?

Solution

We begin by assuming fcc;

d =4M

83/2NAr3=

4 x (107.9 g mol-1)

83/2 x (6.022 x 1023 mol-1) x (1.44 x 10-8 cm)3

= 10.6 g cm-3. Hence silver has fcc structure.


5.14 Ionic Structures5.14 Ionic Structures

- Ionic structures are (like metallic structures) close packed, with anions forming a slightly expanded close-packed structure with smaller cations occupying some of the enlarged holes in the expanded lattice.

- Smaller tetrahedral hole for small cations,larger octahedral hole for somewhat bigger cations

Coordination number: the number of ions of opposite charge immediately surrounding a specific ion

Radius ratio ()


Rock-salt structure of NaCl

Has Cl– ions forming an fcc structure (expanded ccp) and Na+ ions occupying the octahedral hole

Found for a number of minerals having ions of the same charge number, including NaCl, KBr, RbI, MgO, CaO, AgCl


- Rock-salt (fcc) structure: (6,6) coordination, the coordination numbers of the cations and the anions are both 6.

- Common whenever the cations and the anions have very different radii; the cations can fit into the octahedral holes in a fcc array of anions; 0.4 < < 0.7


Cesium chloride structure

- Cl– ions form an expanded primitive cubic array and Cs+ ions occupy large cubic holes (bcc).

- The radii of the cations and anions are similar with > 0.7.

r(Cs+) = 167 pm, r(Cl-) = 181 pm; = 0.923

- It has (8,8) coordination and is less common.Examples: CsBr, CsI, TlCl, TlBr


Zinc-blend (sphalerite) structure, < 0.4

- NiAs: strong Ni–As covalent characterhcp As with Ni in all octahedral holes(6,6)-coordination

- ZnS: expanded ccp array of S2– and small Zn2+ in half of the tetrahedral holes: (4,4)-coordination


Self-Test 5.6A

Predict (a) the likely structure and (b) the coordinationtype of ammonium chloride. Assume that theammonium ion can be approximated as a sphere with aradius of 151 pm.

Solution

Radius ratio, =Radius of smaller ion

Radius of larger ion

=151 pm

181 pm= 0.834

This indicates (a) a cesium chloride structure with (b) (8,8)-coordination.


Self-Test 5.7B

Estimate the density of cesium iodide from its crystalstructure.

Solution

r(Cs+) = 170 pm; r(I-) = 220 pm

Length of diagonal, b = 170 + 2(220) + 170 pm = 780 pm

Length of side a = 780/3 = 450 pm12

CsI has a cesium chloride (bcc) type structure.

Hence unit cell volume is = 9.11 x 107 pm3 or 9.11 x 10-23 cm3

(1 pm3 = 10-30 cm3)

Each bbc unit cell has one Cs+ ion and one I- ion,Density = mass/volume =

(132.91 + 126.90) g mol-1

(6.022 x 1023 mol-1)9.11 x 10-23 cm3 = 4.74 g cm-3


THE IMPACT ON MATERIALS(Sections 5.15-5.16)

5.15 Liquid Crystals5.15 Liquid Crystals Liquid crystals are substances that flow like viscous liquids, but their molecules lie in a moderately orderly array.

- mesophase: an intermediate state of matter with the fluidity ofa liquid and some of the molecular order of a solid

- responsive to changes in temperature and electric fields

- isotropic vs. anisotropic (due to ordering of rodlike molecules)

- p-azoxyanisole, a long and rodlike liquid crystal molecule


Three classes of liquid crystals according to structure

- Nematic phase: the molecules lie together, all in the same directionbut staggered.

- Smectic phase: the molecules line up like soldiers on parade and form layers.

- Cholesteric phase: the molecules form ordered layers, but neighboringlayers have molecules at different angles and so the liquid crystal has a helical arrangement of molecules.


- Lyotropic LC: ordering effects induced by a solvent. The molecules are

amphiphiles (surfactants), with hydrophilic and hydrophobic parts in one molecule

Uses: watches, LCD, thermometers, ···

- Thermotropic LC: made by melting solids. They have long rod shaped Molecules. Example: p-azoxyanisole

Example: sodium lauryl sulfate

Two classes of liquid crystals according to method of preparation


- LCD television or monitor

Layers of a liquid crystal in a nematic phase lie between the surfaces of twoglass or plastic plates.


5.16 Ionic Liquids5.16 Ionic Liquids

- A liquid at room temperature is likely to be a molecular substance.There will be nonionic, weak intermolecular interactions and it will have a relatively high vapor pressure.

Ionic liquids: These are characterized by a relatively small anions (BF4

–) + large organic cation (e.g. 1-butyl-3-methylimidazolium ion), preventing crystallization.

Low vapor pressure, novel solvent properties: reducing pollution


X-RAY DIFFRACTION

The Technique

- interference: When two or more waves pass through the same region,interference is observed as an increase (constructive) ora decrease (destructive) in the total amplitude of the wave.

- diffraction: interference between waves that arises when there is anobject in their path

- Regular layers of atoms in a crystal giving a diffraction pattern


- Why x-rays?

The separation between layers of atoms in a crystal ~ 100 pm

corresponding to the x-ray region

Experimental Techniques

- Powder diffraction technique: a monochromatic (single-frequency)beam of x-rays is directed at a powdered sample spread on asupport.

- Bragg equation 2d sin =with the angles , to the spacing d, for x-rays of wavelength


- Single-crystal diffraction technique

1) Growing a perfect single crystal of the sample (~ 0.1 mm)

- very challenging!

2) Placing the crystal at the center of a four-circle diffractometer

3) Fourier synthesis (conversion) into the locations of the atoms

- raw data including intensities and angles of the diffraction

- description of the atomic locations, bond lengths, and angles

big anglediffraction

Narrowspacing of10 bases

per turn ofhelix

X patternof DNA

Helix witha regularpitch and

radius

Documents

2012 General Chemistry I 1 Chapter 5. LIQUIDS AND SOLIDS 2012 General Chemistry I INTERMOLECULAR FORCES LIQUID STRUCTURE 5.1 The Origin of Intermolecular