Introduction to Polymers, Lecture notes of Chemistry

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A BRIEF INTRODUCTION TO POLYMERIC MATERIALS

by: Prof. Mark E. Tuttle Dept Mechanical Engineering M/S 352600 University of Washington Seattle, WA 98195-

1.0 Polymeric Materials.

A structural engineer who wishes to use adhesives must understand at least the rudiments of polymer chemistry, in much the same way that a structural engineer working with metal alloys must understand at least the rudiments of metallurgy. A brief introduction to polymer material science is given in the following subsections. The underlying objective of this discussion is to simply define some of the terms commonly used to classify or otherwise describe polymers. Since this introduction is necessarily brief, the reader desiring a more detailed discussion is referred to any of the many excellent introductory texts devoted to organic chemistry and/or polymeric materials, such as references 1 through 3. Although very "dated", you may also find the brief tutorial article by Richardson and Kierstead (Ref 4) helpful.

1.1 Atomic Structure: All matter is composed of atoms. An atom consists of a small and very dense core called the nucleus , which is "orbited" by smaller particles called electrons. The nucleus is made up of both protons and neutrons. The mass of protons and neutrons is about the same (1.6726 X 10-24 g and 1.6750 X 10-24 g, respectively), whereas the mass of electrons is four orders of magnitude lower (approximately 9.11 X 10-28 g). The proton is positively charged, the electron is negatively charged, and the neutron is electrically neutral. The magnitude of the electrical charge associated with protons and electrons is identical (approximately 1.602 X 10-19 Coulombs). The electrical charge of these subatomic particles is usually described in relative terms. That is, instead of specifying the charge in "Coulombs," the proton is said to have a charge of "+1", while the electron has a charge of "-1."

The atomic numbe r of an atom equals the number of protons present within the nucleus, and is commonly designated by the symbol "Z." The number of neutrons within the nucleus is called

"break" the bond. On the other hand, electron(s) in the outermost shell are less tightly bonded to the nucleus. In general, much less energy is required to "break" the bond between these outer electrons and the nucleus and, depending on the atom, the bond between an outer electron and the nucleus may be broken or formed relatively easily. Almost all chemical reactions involve the electrons within the outermost shells, and the electrons within these outermost shells are called valence electrons.

Since an atom may "lose" or "gain" valence electron(s), the atom may develop a net electrical charge. An atom that has lost or gained an electron is called an ion. If the atom loses an electron it becomes positively charged and is called a cation. Conversely, if the atom gains an electron it becomes negatively charged and is called an anion.

A periodic table of the elements is shown in Figure 1^1. A total of 109 elements have been

identified. The symbol used to designate each element as well as it’s atomic number and atomic

mass number are shown. One additional item of interest shown in Figure 1 is the concept of a

group of elements, represented by several of the vertical columns within the periodic table. Group

numbers are indicated in Figure 1 by Roman numerals ranging from I to VII. All of the elements

within a group have similar properties because the electron configurations in their outermost

electron shell are similar. That is, the group number of an element equals the total number valence

electrons within the outermost shells of the atom.

(^1) This figure is based on a similar table shown in Ref [1], and does not contain all of the information that ordinarily appears in a periodic table of the elements.

Ni

Ac^

89

H^

1 1.0079Li^

3 6.941Na^

11 22.9898K^

19 39.098Rb 37 85.468Cs^

55 132.905Fr

(^87) (223)

Be9.01218Mg24.305Ca^

20 40.08Sr^

38 87.62Ba^

56 137.33Ra

(^4 88) 226.

Sc^

21 44.9559Y^

39 88.906La^

57

Ti^

22 47.88Zr^

40 91.22Hf^

72 178.49Unq 104 (261)

V^

23 50.9415Nb

41 92.906Ta^

73 180.948Unp

105 (262)

Cr^

24 51.966Mo

42 92.94W^

74 183.85Unh

106 (263)

Mn

25 54.938Tc^

43 (98) Re^

75 186.207Uns

107 (262)

Fe^

26 55.847Ru 44 101.07Os^

76 190.2Uno

108 (265)

Co^

27

28

Cu^

29

Zn^

30

Ga^

31

Ge^

32

As^

33

Se^

34

Br^

35

Kr^

36

Rh^

45

Pd^

46

Ag^

47

Cd^

48

In^

49

Sn^

50

Sb^

51

Te^

52

I^

53

Xe^

54

Ir^

77

Pt^

78

Au^

79

Hg 80

Tl^

81

Pb^

82

Bi^

83

Po^

84 (209)

At^

85 (210)

Rn^

86 (222)

Ce^

58 140.12Th^

90

Pr^

59 140.907Pa

(^91) 231.

Nd 60 144.24U^

92

Pm 61 (145)Np 93

Sm

62 150.36Pu^

94 (244)

Eu^

63

Gd^

64

Tb^

65

Dy 66

Ho 67

Er^

68

Tm

69

Yb 70

Lu^

71

Am

(^95) (243)

Cm

96 (247)

Bk^

97 (247)

Cf^

98 (251)

Es^

99 (252)

Fm

(^100) (257)

Md

(^101) (258)

No^

102 (259)

Lr^

103 (260)

Al^

13

Si^

14

P^

15

S^

16

Cl^

17

Ar^

18

B^

5

C^

6

N^

7

O^

8

F^

9

Ne^

10 He 20. 2

Une

109 (266)

H^

1 ElementSymbol 1.

AtomicNumber Atomic mass number(numbers in parathesesindicates the atomicmass number of isotopewith longest half-life)

I

II^

III

IV

V^

VI

VII

GroupNumbers 12

Figure 1(a): Periodic Table of the Elements (following [1])

1.2 Chemical Bonds: The term "chemical bond" refers to the attractive forces that cause two (or

more) atoms to bond together so as to form a recognizable chemical entity. The "new" chemical

entity normally exhibits properties that differ from the original constituent atom(s). There are three

fundamental types of chemical bonds: metallic bonds , ionic bonds , and covalent bonds. Note that

these forces occur at the atomic level. A second category of forces can be defined at the molecular

level. The magnitude of these intermolecular forces (or "secondary forces" ) are much less than

those associated with chemical bonding, as will be discussed in section 1.3.

1.2.1. Metallic Bonds. Metallic bonding is illustrated schematically in Figure 2(a). Metallic bonding generally occurs for elements that have only one or two valence electrons, since these elements can easily lose electrons to form positively charged cations. Each atom contributes an electron(s) to a "sea" of electrons surrounding the cations. That is, the electrons do not "belong" to any individual cation but rather are free to move within the atomic structure. Metallic bonding is most commonly encountered in the elemental metals or metallic alloys. The very high electrical and thermal conductivities exhibited these materials is a direct result of the mobile electron sea which exists at the atomic level. Note that independent molecules do not exist within substances formed by metallic bonding. Most polymers do not involve metallic bonds, and hence the polymer chemist is not often concerned with this form of chemical bonding.

1.2.2. Ionic Bonds. The ionic bond is illustrated in Figure 2(b). Ionic bonding occurs as a result of the electrostatic attraction between positive cations and negative anions. Ionic compounds are substances formed by ionic bonding. As in the case of metallic bonding, independent molecules do not exist within ionic compounds. Ionic compounds do not contain mobile electrons, and consequently are poor conductors of electrical or thermal energy.

(+) (+) (+) (+) (+) (+) (+) (+)^ (+)

e(-)^ e(-)

e(-)^ e(-)

e(-)

e(-)

e(-)

e(-)

e(-) e(-) e(-)

e(-)

cations

free electrons electron "sea"

(a) Illustration of Metallic Bonding

(+)

(+) (-) (+)

(-) (-)

(+) (-) (+)

cations

anion

(b) Illustration of Ionic Bonding

shared pair of electrons

nucleus

(c) Illustration of Covalent Bonding Figure 2: Schematic Representation of the Three Fundamental Types of Chemical Bonds: Metallic, Ionic, and Covalent Bonding [1]

The preceding discussion has implied that the electrons within a covalent bond are shared "equally" by both atomic nuclei. This is only true when the two atomic nuclei that form the bond attract the electrons pair equally. Since the magnitude of attractive force is established by the number of protons in the nucleus, the electron pair is only shared equally if both atoms have the same number of protons, i.e., if both atoms are of the same element. Examples of this type of molecule are H 2 , Cl 2 , or N 2. In these cases the electron density is equal for both atoms within the

molecule and the bond is called a nonpolar covalent bond.

In contrast, if a covalent bond is formed in which one atom exerts a stronger attractive force than the other, then the electron density around one atom is greater than the other and a polar covalent bond is developed. In effect, the atom with the greater electron density develops a partial

negative charge ( δ −, say) while the atom with the lower electron density develops a partial positive

charge ( δ +, say). This usually occurs when a covalent bond is formed between atoms of different

elements (because the nucleus of the two elements possess a different number of protons), but can also occur between atoms of the same element if the atoms are a part of a larger molecule. The difference between a nonpolar covalent bond and a polar covalent bond is shown schematically in Figure 3.

A polar molecule exhibits non-uniform electron densities, and can be treated as a dipole. That is, although the molecule remains electrically neutral as a whole, one region of the molecule develops a negative charge while a second region develops a positive charge of equal magnitude. The two electrically-charged regions exist at a specific distance from each other, forming an electrical dipole. The negatively-charged region of a polar molecule is attracted to the positively-

δ+^ δ-

(a) Nonpolar Covalent Bond (b) Polar Covalent Bond

Figure 3: Difference Between a "Nonpolar" and "Polar" Covalent Bond

δ-^ δ+ δ-^ δ+ δ-^ δ+

δ+^ δ-^ δ+^ δ-^ δ+^ δ-

δ+^ δ-^ δ+^ δ-^ δ+^ δ-

Figure 4: Schematic Representation of Dipole-Dipole Interactions in a Polar Covalent Compound [1]

negatively-charged electron clouds. Hence, for any covalent bond there is an "equilibrium spacing" between nuclei; that is, the distance between nuclei at which the forces of attraction are exactly balanced by the forces of repulsion. This equilibrium spacing between nuclei is called the covalent bond length. The energy needed to pull the two nuclei apart, thereby destroying the covalent bond, is called the bond energy. 3 Bond lengths and bond energies for covalent bonds often encountered in polymeric substances are listed in Table 1. Note that the bond lengths are exceedingly small; on the order of a tenth of a nanometer (1 nm = 10-9 m). Also note that bond energies are often reported in units of either kcal/mole or kJ/mole, where 1 calorie = 4.19 Joules.

1.3 Intermolecular Forces: The phrase "intermolecular forces" refers to forces of attraction (or repulsion) between molecules or sections of molecules, and are generally of much lower magnitude than interatomic forces associated with chemical bonding. Although names and classifications vary from author-to-author, four principal types of intermolecular forces are generally recognized: dipole-dipole forces , dipole-induced dipole forces , London forces (also called dispersion or van der Waal forces) , and hydrogen bonding. Intermolecular forces are often referred to as "secondary bonds." 4

1.3.1: Dipole-dipole forces: The source of dipole-dipole forces has been described in section 1.2.3 - they arise because the positive end of a polar molecule is attracted to the negative end of a neighboring polar molecule. These forces are sometimes referred to as permanent dipole- dipole forces , because they arise from polar molecules that are "permanent" dipoles.

1.3.2: Dipole-induced dipole forces: A polar molecule (i.e., a permanent dipole) may cause a shift in the electron density of a neighboring nonpolar molecule. Thus, the initially nonpolar molecule becomes a temporary or induced dipole. Dipole-induced dipole forces (sometimes

(^34) Various authors also refer to the bond energy as the "bond strength" or the bond "dissociation energy". Some authors do not include hydrogen bonds within the category of "secondary bonds."

referred to as Keesom forces ) are the forces of attraction between the permanent and induced dipoles.

1.3.3: London (or dispersion) forces: London forces are also due to shifts in the electron density of a nonpolar molecule, which result in temporary dipoles. However, in contrast to Keesom forces, London forces are caused by interactions between two temporary dipoles. One end (or region) of a molecule momentarily develops a slight positive charge, say, which consequently means that the other end (or a second region) momentarily develops a slight negative charge. The positive end/region of the molecule will induce a slight negative charge in an adjacent region of a neighboring molecule, causing an asymmetric electron density in this neighboring molecule. London forces are the forces of attraction between fluctuating induced dipoles. London forces are extremely significant in that they are present in all covalently bonded substances, whether polar or non-polar.

1.3.4: Hydrogen bonding: Although hydrogen bonding is customarily listed as a separate type of intermolecular force, in reality hydrogen bonding is simply an unusually strong dipole- dipole force. It occurs when a hydrogen atom is covalently bonded to a relatively small but strongly electronegative atom. In particular, strong hydrogen bond forces result when a hydrogen atom is covalently bonded to a fluorine, nitrogen, or oxygen atom. Since the hydrogen atom carries only a single electron and proton, once the covalent bond is formed with these larger atoms the much smaller hydrogen atom has very little electron density surrounding it, and a strongly polar molecule results. The hydrogen atom develops an unusually large (partial) positive charge. Therefore, in addition to being covalently bonded to an electronegative atom, the hydrogen atom is strongly attracted to a second surrounding electronegative atom.

1.3.5 The strength of intermolecular forces: Inter-molecular forces are present in all molecular substances, and determine whether the substance is a gas, a liquid, or a solid at a given

1.4 Fundamental Aspects of Polymer Molecules: The term "polymer" comes from the Greek words poly (meaning "many") and mers (meaning "units"). At the molecular level polymers consist of extremely long, chain-like molecules. Polymer molecules are typically made up of thousands of repeating chemical units, and have molecular weights ranging from about 10^3 to 10^7.

As an illustrative example, consider the single chemical mer shown in Figure 5. This mer is called ethylene (or ethene ), and consists of two carbon atoms and four hydrogen atoms. The two lines between the carbon (C) atoms indicate a double covalent bond whereas the single line between the hydrogen (H) and carbon atoms represents a single covalent bond. The chemical composition of the ethylene mer is written C 2 H 4 or CH 2 =CH 2. Under the proper conditions one of the double

covalent bonds between the two carbon atoms can be broken, which allows each of the two carbon atoms to form a new covalent bond with a carbon atom in a neighboring mer. In this way three ethylene mers form a "new" molecule, whose atomic weight is three times as great as the initial mer. If "n" ethylene mers join together, the chemical composition of the resulting molecule can be represented C2nH4n, where n is any positive integer. In this way a "chain" of ethylene mers join

together to form the well-known polymer polyethylene , as shown in Figure 6. A typical polyethylene molecule may contain 50,000 carbon atoms or more. The process of causing a monomer to chemically react and form long molecules in this fashion is called polymerization , and the number of repeating units which make up the molecule is called the degree of polymerization.

The single ethylene unit is an example of a monomer. At room temperatures a bulk sample of the ethylene monomer is a low-viscosity fluid. If two ethylene monomers bond together the resulting chemical entity has two repeating units and is called a " dimer ." Similarly, the chemical entity formed by three repeating units is called a " trimer ." The molecular weight of a dimer is twice that of the monomer, the molecular weight of a trimer is three times that of the monomer, etc. Prior

H

H

C C

H

H

Figure 5: The monomer "ethylene"

H

H

C C

H

H (^) n

C

H

H

C

H

H

( groupend ) (^ groupend )

ethylene "repeat" unit Figure 6: The polymer "polyethylene"

H

H

H

C 6

N

H N

H H

n

+ H O^ O

H

H

C 4

C

O C

O H

n

{hexamethylene diamine} {adipic acid}

H

H

H

C 6

N

H N

H

n

H

H

C 4

C

O C

O

O H + (2n-1)

H

H

O

{hexamethylene adipamide or "nylon 66"} (^) {water}

Figure 7: The Polymer "Nylon 66"

A monofunctional mer can form only one covalent bond, and therefore cannot exist as the repeat unit in a polymer, although a monofunctional mer can exist within a polymer as an end group. Trifunctional or tetrafunctional mers can form three or four covalent bonds, respectively.

Although in the case of polyethylene the repeat unit is equivalent to the original ethylene monomer, this is not always the case. In fact, in many instances the repeat unit is derived from two (or more) monomers. A typical example is Nylon 66. The polymerization process for this polymer is shown schematically in Figure 7. Two monomers are used to produce Nylon 66: hexamethylene diamine (chemical composition: C 6 H 16 N 2 ) and adipic acid (chemical composition: COOH(CH 2 ) 4 COOH). Note that the repeat unit of Nylon 66 (hexamethylene adipamide) is not

equivalent to either of the two original monomers. A low-molecular weight byproduct (i.e., water) is produced during the polymerization of Nylon 66. This is a characteristic of condensation polymers. That is, if both a high-molecular weight polymer as well as a low-molecular weight byproduct is formed during the polymerization process, the polymer is classified as a condensation polymer. Conversely, addition polymers are those for which no byproduct is formed during the polymerization process, which implies that all atoms present in the original monomer(s) occur somewhere within the repeat unit. Generally speaking, condensation polymers shrink to a greater extent during the polymerization process than do addition polymers. Residual stresses caused by shrinkage during polymerization (sometimes referred to as "cure stresses") are often a concern in adhesive bonding, and hence difficulties with residual stresses can be minimized if an addition polymer is used in these structural applications.

1.3 Covalent Bond Angles: As previously discussed an individual molecule consists of elemental atoms bonded together via covalent bonds. A particularly simple molecule is the gas methane, CH 4. In this case the carbon atom is bonded to four hydrogen atoms via four single covalent bonds. A 3- D sketch of a single methane molecule is shown in Figure 8. As indicated, if the single molecule

were completely isolated from all other “external effects” (where an “external effect” might be another methane molecule, for example), then the molecule will take on the shape of a four-sided regular tetrahedron. A covalent bond angle is defined as the angle between three neighboring atoms in a molecule. For example, the bond angle formed by the H-C-H atoms in a methane molecule (isolated from all external effects) is 109° 28’, as shown in Figure 8. This bond angle can be considered to the “equilibrium” value. That is, if external forces are present (such as the close proximity of another methane molecule, for example), then the molecule may be distorted, i.e., the covalent bond angle may deviate from 109° 28’. Note that since some work must be done on the molecule to cause this distortion, the “internal energy” of the molecule has been increased due to this distortion.

C H

H

H H

109 28'o

Figure 8: 3-D sketch of methane molecule, isolated from all external effects

Of course, multiple covalent bond angles can be defined for more complex molecules, since complex molecules may involve many different elemental atoms and/or greater numbers of atoms than the simple methane molecule. Nevertheless, it is possible to calculate “equilibrium” values for all bond angles involved in a molecule, and if the bond angles deviate from their equilibrium values the “internal energy” of the molecule has been increased.

A second form of molecular “distortion” occurs through rotation about covalent bonds. This is illustrated in Figure 9 for the ethane molecule (C 2 H 6 ). A 3-D sketch of the molecule is shown in