Interatomic Bonding

12 Key excerpts on "Interatomic Bonding"

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  eBook - PDF

  Fundamentals of Materials Science and Engineering

  An Integrated Approach

  • William D. Callister, Jr., David G. Rethwisch(Authors)
  • 2021(Publication Date)
  • Wiley

    (Publisher)

  Stewart/Science Source • 19 WHY STUDY Atomic Structure and Interatomic Bonding? An important reason to have an understanding of Interatomic Bonding in solids is that in some instances, the type of bond allows us to explain a material’s properties. For example, consider carbon, which may exist as both graphite and diamond. Whereas graphite is relatively soft and has a “greasy” feel to it, diamond is the hardest known material. In addition, the electrical properties of diamond and graphite are dissimilar: diamond is a poor conductor of electricity, but graphite is a reasonably good conductor. These disparities in properties are directly attributable to a type of intera- tomic bonding found in graphite that does not exist in diamond (see Section 3.9). Learning Objectives After studying this chapter, you should be able to do the following: 1. Name the two atomic models cited, and note the differences between them. 2. Describe the important quantum-mechanical principle that relates to electron energies. 3. (a) Schematically plot attractive, repulsive, and net energies versus interatomic separation for two atoms or ions. (b) Note on this plot the equilibrium separation and the bonding energy. 4. (a) Briefly describe ionic, covalent, metallic, hydrogen, and van der Waals bonds. (b) Note which materials exhibit each of these bonding types. Some of the important properties of solid materials depend on geometric atomic ar- rangements and also the interactions that exist among constituent atoms or molecules. This chapter, by way of preparation for subsequent discussions, considers several fun- damental and important concepts—namely, atomic structure, electron configurations in atoms and the periodic table, and the various types of primary and secondary inter- atomic bonds that hold together the atoms that compose a solid. These topics are re- viewed briefly, under the assumption that some of the material is familiar to the reader.

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  Introduction to the Physics and Chemistry of Materials

  • Robert J. Naumann(Author)
  • 2008(Publication Date)
  • CRC Press

    (Publisher)

  3 Chemical Bonding The ability of atoms and molecules to form chemical bonds is the de fi ning feature of the structure and properties of solids. The types of bonds that are formed determine if the material will be a metal, a ceramic, or a polymer, and whether the material will conduct electricity, transmit light, or be magnetic. 3.1 What Holds Stuff Together? All matter that we deal with on an everyday basis is held together by electrical forces that form chemical bonds. These forces are manifested in different ways, depending upon which elements are involved. There are three type of primary bonds: (1) the metallic bond in which electrons become detached from atoms when they come together so the ion cores become mutually attracted to the sea of electrons surrounding them; (2) the covalent bond in which atoms become mutually attracted by sharing electrons in order to form closed electron shells; and (3) the ionic bond in which a mutual attraction occurs when one or more electrons leaves a metal atom to complete an atomic shell of a nonmetallic atom forming an oppositely charged ion pair. Much weaker bonds, such as the hydrogen bond, which arise from dipolar attractions between molecules when a hydrogen atom becomes covalently bonded to an O, N, or F atom, or to the van der Waals bond, which arises from induced dipole – dipole interactions, play a secondary role in the structure of materials. Understanding these basic forces that hold materials together is crucial to understanding the structure and properties of materials. We shall start with the ionic bond since conceptu-ally it is the easiest to visualize and it lends itself to a simple analytical model. 3.2 Ionic Bonding The ionic bond is the strongest chemical bond, ranging from 10.5 eV for LiF to 5.8 eV for CsI, but it can only act between two (or more) dissimilar atoms.

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  Fundamentals of Materials Science and Engineering

  An Integrated Approach

  • William D. Callister, Jr., David G. Rethwisch(Authors)
  • 2022(Publication Date)
  • Wiley

    (Publisher)

  Courtesy of Jeffrey Karp, Robert Langer and Alex Galakatos Courtesy of Jeffrey Karp, Robert Langer and Alex Galakatos Barbara Peacock/Getty Images Paul D. Stewart/Science Source Atomic Structure and Interatomic Bonding 2 2.2 Fundamental Concepts  21 2.1 | | INTRODUCTION 1 Protons, neutrons, and electrons are composed of other subatomic particles such as quarks, neutrinos, and bosons. However, this discussion is concerned only with protons, neutrons, and electrons. WHY STUDY Atomic Structure and Interatomic Bonding? An important reason to have an understand- ing of Interatomic Bonding in solids is that in some instances, the type of bond allows us to explain a material’s properties. For example, consider carbon, which may exist as both graph- ite and diamond. Whereas graphite is relatively soft and has a “greasy” feel to it, diamond is the hardest known material. In addition, the electrical properties of diamond and graphite are dissimilar: diamond is a poor conductor of electricity, but graphite is a reasonably good conductor. These disparities in properties are directly attributable to a type of Interatomic Bonding found in graphite that does not exist in diamond (see Section 3.9). LEARNING OBJECTIVES After studying this chapter, you should be able to do the following: 1. Name the two atomic models cited, and note the differences between them. 2. Describe the important quantum-mechanical principle that relates to electron energies. 3. (a) Schematically plot attractive, repulsive, and net energies versus interatomic separation for two atoms or ions. (b) Note on this plot the equilibrium separation and the bonding energy. 4. (a) Briefly describe ionic, covalent, metallic, hydrogen, and van der Waals bonds. (b) Note which materials exhibit each of these bonding types.

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  Materials Science and Engineering, P-eBK

  An Introduction

  • William D. Callister, Jr., David G. Rethwisch, Aaron Blicblau, Kiara Bruggeman, Michael Cortie, John Long, Judy Hart, Ross Marceau, Ryan Mitchell, Reza Parvizi, David Rubin De Celis Leal, Steven Babaniaris, Subrat Das, Thomas Dorin, Ajay Mahato, Julius Orwa(Authors)
  • 2020(Publication Date)
  • Wiley

    (Publisher)

  ionic bond A coulombic interatomic bond that exists between two adjacent and oppositely charged ions. isotope Atoms of the same element that have different atomic masses. metallic bond A primary interatomic bond involving the nondirectional sharing of nonlocalised valence electrons (‘sea of electrons’) that are mutually shared by all the atoms in the metallic solid. mole The quantity of a substance corresponding to atoms or molecules. Pauli exclusion principle The postulate that for an individual atom, at most two electrons, which necessarily have opposite spins, can occupy the same state. periodic table The arrangement of the chemical elements with increasing atomic number according to the periodic variation in electron structure. Nonmetallic elements are positioned at the far right‐hand side of the table. polar molecule A molecule in which there exists a permanent electric dipole moment by virtue of the asymmetrical distribution of positively and negatively charged regions. primary Interatomic bonds that are relatively strong and for which bonding energies are relatively large. Primary bonding types are ionic, covalent, and metallic. quantum mechanics A branch of physics that deals with atomic and subatomic systems; it allows only discrete values of energy. By contrast, for classical mechanics, continuous energy values are permissible. quantum number A set of four numbers, the values of which are used to label possible electron states. Three of the quantum numbers are integers that specify the size, shape, and spatial orientation of an electron’s probability density; the fourth number designates spin orientation. secondary bond Interatomic and intermolecular bonds that are relatively weak and for which bonding energies are relatively small. Normally, atomic or molecular dipoles are involved. Examples of secondary bonding types are van der Waals forces and hydrogen bonding.

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  Metals and Materials

  Science, Processes, Applications

  • R. E. Smallman, R J Bishop(Authors)
  • 2013(Publication Date)
  • Butterworth-Heinemann

    (Publisher)

  <3000 kg nr 3 . 1.4 Interatomic Bonding in materials Matter can exist in three states and as atoms change directly from either the gaseous state (desublima-tion) or the liquid state (solidification) to the denser solid state, the atoms form aggregates in three-dimensional space. Bonding forces develop as atoms are brought into proximity to each other. Sometimes these forces are spatially-directed. The nature of the bonding forces has a direct effect upon the type of solid structure which develops and therefore upon the physical properties of the material. Melting point provides a useful indication of the amount of thermal energy needed to sever these interatomic (or interionic) bonds. Thus, some solids melt at relatively low temperatures (m.p. of tin = 232°C) whereas many ceramics melt at extremely high temperatures (m.p. of alumina exceeds 2000°C). It is immediately apparent that bond strength has far-reaching implications in all fields of engineering. Customarily we identify four principal types of bonding in materials, namely, metallic bonding, ionic bonding, covalent bonding and the compara-tively much weaker van der Waals bonding. However, in many solid materials it is possible for bonding to be mixed, or even intermediate, in character. We will first consider the general chemi-cal features of each type of bonding; in Chapter 2 we will examine the resultant disposition of the assembled atoms (ions) in three-dimensional space. •Atomic mass is now expressed relative to the datum value for carbon (12.01). Thus, a copper atom has 63.55/12.01 or 5.29 times more mass than a carbon atom. 8 Metals and Materials (a) Sodium (Z = 11) llll@llllll©llllll©llll (b) Magnesium (Z = 12) and oxygen (Z = 8) !cf f (c) Carbon (Z = 6) (d) Polarized atoms Figure 1.3 Schematic representation of (a) metallic bonding, (b) ionic bonding, (c) covalent bonding and (d) van der Waals bonding.

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  eBook - PDF

  Materials Science and Engineering

  An Introduction

  • William D. Callister, Jr., David G. Rethwisch(Authors)
  • 2018(Publication Date)
  • Wiley

    (Publisher)

  • Bonding force and bonding energy are related to one another according to Equa- tions 2.5a and 2.5b. • Attractive, repulsive, and net energies for two atoms or ions depend on interatomic separation per the schematic plot of Figure 2.10b. • For ionic bonds, electrically charged ions are formed by the transference of valence electrons from one atom type to another. • There is a sharing of valence electrons between adjacent atoms when bonding is covalent. • Electron orbitals for some covalent bonds may overlap or hybridize. Hybridization of s and p orbitals to form sp 3 and sp 2 orbitals in carbon was discussed. Configurations of these hybrid orbitals were also noted. • With metallic bonding, the valence electrons form a “sea of electrons” that is uni- formly dispersed around the metal ion cores and acts as a form of glue for them. • Relatively weak van der Waals bonds result from attractive forces between electric dipoles, which may be induced or permanent. • For hydrogen bonding, highly polar molecules form when hydrogen covalently bonds to a nonmetallic element such as fluorine. • In addition to van der Waals bonding and the three primary bonding types, covalent– ionic, covalent–metallic, and metallic–ionic mixed bonds exist. • The percent ionic character (%IC) of a bond between two elements (A and B) depends on their electronegativities (X’s) according to Equation 2.16. • Correlations between bonding type and material class were noted: Polymers—covalent Metals—metallic Ceramics—ionic/mixed ionic–covalent Molecular solids—van der Waals Semi-metals—mixed covalent–metallic Intermetallics—mixed metallic–ionic Electrons in Atoms The Periodic Table Bonding Forces and Energies Primary Interatomic Bonds Secondary Bonding or van der Waals Bonding Mixed Bonding Bonding Type- Material Classification Correlations

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  Fundamentals of Materials Science and Engineering

  An Integrated Approach

  • William D. Callister, Jr., David G. Rethwisch(Authors)
  • 2016(Publication Date)
  • Wiley

    (Publisher)

  Electrons in Atoms The Periodic Table Bonding Forces and Energies Primary Interatomic Bonds Secondary Bonding or van der Waals Bonding 44 • Chapter 2 / Atomic Structure and Interatomic Bonding • In addition to van der Waals bonding and the three primary bonding types, covalent– ionic, covalent–metallic, and metallic–ionic mixed bonds exist. • The percent ionic character (%IC) of a bond between two elements (A and B) de- pends on their electronegativities (X’s) according to Equation 2.16. • Correlations between bonding type and material class were noted: Polymers—covalent Metals—metallic Ceramics—ionic/mixed ionic–covalent Molecular solids—van der Waals Semi-metals—mixed covalent–metallic Intermetallics—mixed metallic–ionic Mixed Bonding Bonding Type- Material Classification Correlations List of Symbols Symbol Meaning A, B, n Material constants E Potential energy between two atoms/ions E A Attractive energy between two atoms/ions E R Repulsive energy between two atoms/ions e Electronic charge ε 0 Permittivity of a vacuum F Force between two atoms/ions r Separation distance between two atoms/ions X A Electronegativity value of the more electronegative element for compound BA X B Electronegativity value of the more electropositive element for compound BA Z 1 , Z 2 Valence values for ions 1 and 2 Equation Summary Equation Page Number Equation Solving For Number 2.5a E = ∫ F dr Potential energy between two atoms 29 2.5b F = dE dr Force between two atoms 29 2.9 E A = − A r Attractive energy between two atoms 30 2.11 E R = B r n Repulsive energy between two atoms 31 2.13 F A = 1 4πε 0 r 2 (  Z 1  e )(  Z 2  e ) Force of attraction between two isolated ions 33 2.16 %IC = { 1 − exp[ −(0.25)( X A − X B ) 2 ] } × 100 Percent ionic character 41

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  Materials for Engineering

  • J Martin(Author)
  • 2006(Publication Date)
  • Woodhead Publishing

    (Publisher)

  Part I Characterization of engineering materials Passage contains an image

  1

  Structure of engineering materials

  1.1 Crystal structure

  Crystal structure refers to the ordering of atoms into different crystalline arrangements. It is the arrangement of these atoms – the strength and directionality of the interatomic bonds – which determines the ultimate strength of the solid. Techniques involving X-ray or electron diffraction are employed to determine crystal structures, and four types of Interatomic Bonding are recognized: van der Waals, covalent, ionic and metallic. The latter three ‘primary’ bonds are limiting cases, however, and a whole range of intermediate bonding situations also exist in solids.

  The van der Waals force is a weak ‘secondary’ bond and it arises as a result of fluctuating charges in an atom. There will be additional forces if atoms or molecules have permanent dipoles as a result of the arrangement of charge inside them. In spite of their low strength, these forces can still be important in some solids; for example it is an important factor in determining the structure of many polymeric solids.

  Many common polymers consist of long molecular carbon chains with strong bonds joining the atoms in the chain, but with the relatively weak van der Waals bonds joining the chains to each other. Polymers with this structure are thermoplastic, i.e. they soften with increasing temperatures and are readily deformed, but on cooling they assume their original low-temperature properties and retain the shape into which they were formed.

  Covalent bonding is most simply exemplified by the molecules of the non-metallic elements hydrogen, carbon, nitrogen, oxygen and fluorine. The essential feature of a covalent bond is the sharing of electrons between atoms, enabling them to attain the stable configuration corresponding to a filled outermost electron shell. Thus, an atom with n electrons in that shell can bond with only 8 – n

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  Nanotechnology

  Understanding Small Systems, Third Edition

  • Ben Rogers, Jesse Adams, Sumita Pennathur(Authors)
  • 2014(Publication Date)
  • CRC Press

    (Publisher)

  These three kinds of bonds are also what hold together the more complex, large-scale crystalline formations of atoms within solids, as well as nanoscale materials. Metals are one exception. A special arrangement of atoms and electrons known as the metallic bond holds solid metal together. We will dis-cuss these bonding types now. 4.2.1 Ionic Bonding When an atom loses or gains extra electrons, it becomes either positively or negatively charged. The sodium atom, Na, tends to give up its highest-energy electron to become the Binding energy Equilibrium separation Potential energy x x FIGURE 4.1 Total potential energy of a pair of atoms as a function of separation distance, x . At large separation distances, attractive forces pull the atoms toward one another, toward the equilib-rium separation distance. At small separations, repulsive forces dominate. Nanomaterials ◾ 91 positively charged ion, Na + ; whereas the chlorine atom, Cl, tends to pick up a spare electron to fill out its electronic configuration, thereby forming Cl − . Negative and positive ions are drawn together (by Coulomb attraction) and we obtain NaCl (table salt). The Coulomb attraction energy, E , that drives ionic bonding is given by E z z e x = 1 2 0 2 4 πε ε (4.1) Here, x is the distance between the ions, ε 0 is the permittivity of free space, and ε is the dielectric constant of the medium between the ions (or the relative permittivity). The magnitude and sign of the two ions are given in terms of e , the elementary charge ( e = 1.602 × 10 − 19 C), multiplied by the valence state of the ions, z 1 and z 2 . In the case of Na + , for example, z = + 1; for Cl − , z = − 1. An ion such as Ca 2 + has z = + 2, and so on. As a quick review, energy is a force applied over a distance. Taking the negative derivative of the energy equation with respect to distance, or − d E /d x , gives us the force, F , between the atoms.

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  Introduction to Nanoscience and Nanotechnology

  • Gabor L. Hornyak, H.F. Tibbals, Joydeep Dutta, John J. Moore(Authors)
  • 2008(Publication Date)
  • CRC Press

    (Publisher)

  Bonding between and among entities that are larger than atoms or molecules, such as between materials with nanoscale proportions, is the sum total of the individual “bonding elements” integrated across the entire volume and surface of the constituents. The sum total energy in these cases may be quite large. 10.1 E LECTROSTATIC I NTERACTIONS Although all bonding is electrostatic, we especially direct our attention to inter-actions in which there is a clear case for physical charge separation between entities that have an extra electron or two or three or are missing electrons. Permanent dipoles are not charged but exhibit charge separation within their structure. The permanent dipole is due to the difference in electronegativity of its atomic constituents. All chemical bonding is due to interactions between electronic charges that arise from manifestations of Coulomb’s law, given below in its force and energy forms: 498 Introduction to Nanoscience and Nanotechnology F = ⎛ ⎝ ⎜ ⎞ ⎠ ⎟ 1 4 1 2 2 pe e o r q q r (10.2) E = ⎛ ⎝ ⎜ ⎞ ⎠ ⎟ 1 4 1 2 pe e o r q q r (10.3) where F is the force experienced between two elementary charges q , e o and e r are the dielectric permittivity of vacuum and medium, respectively, and r is the dis-tance between the two charges. E is the potential energy between two point charges and is inversely proportional to the distance between them. The value of the vacuum permittivity constant e o is equal to 8.854 × 10 –12 C 2 ·J –1 ·m –1 and that of the elementary charge is 1.6022 × 10 –19 C. Ion–ion, ion–dipole, and dipole– dipole interactions are electrostatic interactions between atoms and/or mole-cules that possess relatively permanent electrostatic charges. Charles-Augustin Coulomb is pictured in Figure 10.4. Repulsive Interactions. Lest we place too little value on repulsive interactions, we shall provide a cursory overview of a few kinds of electrostatic repulsive inter-actions that exist in materials.

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  Nanotechnology

  The Whole Story

  • Ben Rogers, Jesse Adams, Sumita Pennathur(Authors)
  • 2013(Publication Date)
  • CRC Press

    (Publisher)

  These three kinds of bonds are also what hold together Binding energy Equilibrium separation Potential energy x x FIGURE 4.1 Total potential energy of a pair of atoms as a function of separation distance, x . At large separation distances, attractive forces pull the atoms toward one another, toward the equilibrium separation distance. At small separations, repulsive forces dominate. Nanomaterials 81 the more complex, large-scale crystalline formations of atoms within sol-ids, as well as nanoscale materials. Metals are one exception. A special arrangement of atoms and electrons known as the metallic bond holds solid metal together. We will discuss these bonding types now. 4.2.1 Ionic Bonding When an atom loses or gains extra electrons, it becomes either posi-tively or negatively charged. The sodium atom, Na, tends to give up its highest energy electron to become the positively charged ion, Na + , whereas the chlorine atom, Cl, tends to pick up a spare electron to fill out its electronic configuration, thereby forming Cl − . Negative and pos-itive ions are drawn together (by Coulomb attraction) and we obtain NaCl (table salt). TABLE 4.1 Main Types of Bonding Bonding Among Atoms and Molecules Within Solids Ionic Oppositely charged atoms are attracted to one another to make molecules. Example: Na + and Cl − make the NaCl molecule. Crystal structure is formed by an array of atoms held together by opposing charges. Example: Magnesium oxide, a crystal formed by O 2– and Mg 2+ molecules. Covalent Two atoms share electrons; their atomic orbitals overlap to make a molecular orbital. Example: Two hydrogen atoms make the H 2 molecule. Crystal structure is formed by an array of atoms sharing electron orbitals. Example: Diamond, made from C atoms each sharing electrons with four other C atoms. van der Waals (three types) Two atoms are attracted to one another by weak electrostatic forces.

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  Introduction to Nanoscience

  • Gabor L. Hornyak, Joydeep Dutta, H.F. Tibbals, Anil Rao(Authors)
  • 2008(Publication Date)
  • CRC Press

    (Publisher)

  Chapters 13 and 14 are allocated to the biological nanoscience division; they focus on natural nanomaterials and biochemical nanoscience, respectively. It is time to refresh our memories once again concerning chapter 2 on the societal implications of nano. As you proceed through this chapter and the rest of the text, always ask the question, “What are the consequences—good, bad, or indifferent— of these materials?” 10.0 B ONDING C ONSIDERATIONS AT THE N ANOSCALE Chemical and physical interactions between atoms and molecules comprise two sides of the same coin. This apparent dichotomy is once again a product of con-venience, generated to feed our instinct to catalog. Two atoms held together by attractive forces form a chemical bond that in turn yields a molecule (from the Latin moles, “small unit of mass” or “mass barrier” based on the Greek molos, “exertion”). Attractive forces between atoms form strong intramolecular bonds as a result of chemical reactions. We have briefly introduced types of intramolecular bonding such as the covalent bond, the ionic bond, and other types like the metallic bond in chapter 5. We now expand our discussion to the intermolecular bond —a type of interaction that exists between two or more molecules. We define physical processes as those that involve no change in the chemical structure of molecules. We define chemical processes as those that do involve changes in the chemical structure of molecules. If we apply these definitions, in the strictest sense, to intermolecular bonds, they should be classified as a physi-cal interaction because the structure (chemical nature) of the precursor mole-cules is not altered (significantly). However, any type of bonding, regardless of how strong or weak, causes perturbations to the integrated molecular orbital of any preexisting molecular system—perturbations that can be construed to be chemical in nature. Although no bonds are made or broken (except in the case of

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