Metallic Bonding

9 Key excerpts on "Metallic Bonding"

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  eBook - ePub

  Understanding Basic Chemistry

  The Learner's Approach

  • Kim Seng Chan, Jeanne Tan(Authors)
  • 2014(Publication Date)
  • WSPC

    (Publisher)

  to gain electrons across the period!

  In a nutshell, remember:

  •Elements across a period experience increasing ease to gain electrons BUT decreasing ease to lose electrons.

  •Elements down a group experience decreasing ease to gain electrons BUT increasing ease to lose electrons.

  3.1 Metallic Bonding

  Chemical bonds are electrostatic forces of attraction (positive charge attracting negative charge) that bind particles together to form matter. When different types of particles interact electrostatically, different types of chemical bonds are formed. There are four different types of conventional chemical bonds, namely, metallic, ionic, covalent, and intermolecular forces.

  Within a metal, atoms partially lose their loosely bound valence electrons. These electrons are mobile and delocalized, not belonging to any one single atom and yet not completely lost from the lattice.

  A metal can thus be viewed as a rigid lattice of positive ions surrounded by a sea of delocalized electrons. What holds the lattice together is the strong Metallic Bonding — the electrostatic attraction between the positive ions and the delocalized valence electrons.

  Metallic bonds are strong and non-directional. Therefore, when a force is applied across a piece of metal, the metal atoms can slide over one another without breaking of the metallic bonds. This accounts for the malleability (can be deformed into different shapes) and ductility (can be drawn into wires) of metals.

  Since Metallic Bonding is the result of the interaction between the delocalized electrons and the positive ions, the strength of a metallic bond depends on:

  •The number

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  Chemistry

  Structure and Dynamics

  • James N. Spencer, George M. Bodner, Lyman H. Rickard(Authors)
  • 2011(Publication Date)
  • Wiley

    (Publisher)

  The force of attraction that would have to be overcome to separate the metal ions can be thought of as the metallic bond that holds the par- ticles together. Metals exist as extended three-dimensional arrays of atoms, which pack so that each atom can touch as many neighboring atoms as possible. When a metal is heated or beaten with a hammer, the planes of atoms that form the structure can slip past one another, which explains why metals are malleable and ductile (see Section 5.1) Each atom in the structure has a limited number of loosely held valence electrons that it can share with its neighbors. Due to the mobility of the electrons, metals can easily transfer kinetic energy from one atom to another. As a result, they are good conductors of heat. Because the valence electrons in a metal are delocalized instead of being strictly associated with a single atom, they are free to move from one atom to another. The electrons that freely move through the metal are said to form conduction bands. If we draw the metal into a thin wire and connect the wire to a source of an electric current, electrons that enter the wire displace electrons that were already present on the atoms closest to the source of the current. Electrons flow through the conduction band until they eventually displace electrons from the other side of the wire. Metals are therefore good conductors of electricity. Metals in their elemental form (Na, Cu, Fe) do not consist of individual atoms. They are composed of a three-dimensional network of positive metal ions surrounded by a sea of electrons. Consider a sample of pure iron. Because it would contain only iron atoms, we write the chemical formula as Fe. 5.13 The Relationship among Ionic, Covalent, and Metallic Bonds There are significant differences in the physical properties of sodium chloride, sodium metal, and chlorine, as shown in Table 5.4. These differences result from significant differences in both the structure and bonding in these substances.

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  eBook - PDF

  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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  eBook - PDF

  Fundamentals of Materials Science and Engineering

  An Integrated Approach

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

    (Publisher)

  In addition, these free electrons act as a “glue” to hold the ion cores together. Bonding energies and melting temperatures for several met- als are listed in Table 2.3. Bonding may be weak or strong; energies range from 62 kJ/mol for mercury to 850 kJ/mol for tungsten. Their respective melting temperatures are −39°C and 3414°C (−39°F and 6177°F). Metallic Bonding is found in the periodic table for Group IA and IIA elements and, in fact, for all elemental metals. Metals are good conductors of both electricity and heat as a consequence of their free electrons (see Sections 12.5, 12.6, and 17.4). Furthermore, in Section 8.5, we note that at room temperature, most metals and their alloys fail in a ductile manner—that is, fracture occurs after the materials have experienced significant degrees of permanent deforma- tion. This behavior is explained in terms of deformation mechanism (Section 8.3), which is implicitly related to the characteristics of the metallic bond. Tutorial Video: Bonding What Is Metallic Bonding? Figure 2.19 Schematic illustra- tion of Metallic Bonding. - - - - - - - - - - - - - - - - - - - - - - + + + + + + + + + + + + + + + + + + + + + Ion core Metal atom Sea of valence electrons - - Concept Check 2.3 Explain why covalently bonded materials are generally less dense than ionically or metallically bonded ones. (The answer is available in WileyPLUS.) Secondary bonds, or van der Waals (physical) bonds, are weak in comparison to the primary or chemical bonds; bonding energies range between about 4 and 30 kJ/mol. Secondary bonding exists between virtually all atoms or molecules, but its presence may be obscured if any of the three primary bonding types is present. Secondary bonding is evidenced for the inert gases, which have stable electron structures. In addition, second- ary (or intermolecular) bonds are possible between atoms or groups of atoms, which themselves are joined together by primary (or intramolecular) ionic or covalent bonds.

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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)

  In addition, these free electrons act as a “glue” to hold the ion cores together. Bonding energies and melting temperatures for several metals are listed in Table 2.3. Bonding may be weak or strong; energies range from 62 kJ/mol for mercury to 850 kJ/mol for tungsten. Their respective melting temperatures are −39°C and 3414°C (−39°F and 6177°F). Metallic Bonding is found in the periodic table for Group IA and IIA elements and, in fact, for all elemental metals. Metals are good conductors of both electricity and heat as a consequence of their free electrons (see Sections 18.5, 18.6, and 19.4). Furthermore, in Section 7.4, we note that at room temperature, most metals and their alloys fail in a ductile manner—that is, fracture occurs after the materials have experienced significant degrees of permanent deformation. This behavior is explained in terms of a deformation mechanism (Section 7.2), which is implicitly related to the characteristics of the metallic bond. Figure 2.19 Schematic illustra- tion of Metallic Bonding. - - - - - - - - - - - - - - - - - - - - - - + + + + + + + + + + + + + + + + + + + + + Ion core Metal atom Sea of valence electrons - - Concept Check 2.3 Explain why covalently bonded materials are generally less dense than ionically or metallically bonded ones. [The answer may be found in all digital versions of the text and/or at www.wiley.com/college/callister (Student Companion Site).] Secondary bonds, or van der Waals (physical) bonds, are weak in comparison to the primary or chemical bonds; bonding energies range between about 4 and 30 kJ/mol. Secondary bonding exists between virtually all atoms or molecules, but its presence may be obscured if any of the three primary bonding types is present. Secondary bonding is evidenced for the inert gases, which have stable electron structures.

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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)

  For each type, the bonding necessarily involves the valence electrons; furthermore, the nature of the bond depends on the electron structures of the constitu- ent atoms. In general, each of these three types of bonding arises from the tendency of the atoms to assume stable electron structures, like those of the inert gases, by com- pletely filling the outermost electron shell. Secondary or physical forces and energies are also found in many solid materials; they are weaker than the primary ones but nonetheless influence the physical properties of some materials. The sections that follow explain the several kinds of primary and secondary interatomic bonds. primary bond Ionic Bonding Ionic bonding is perhaps the easiest to describe and visualize. It is always found in compounds composed of both metallic and nonmetallic elements, elements situated at the horizontal extremities of the periodic table. Atoms of a metallic element easily give up their valence electrons to the nonmetallic atoms. In the process, all the atoms acquire stable or inert gas configurations (i.e., completely filled orbital shells) and, in addition, an electrical charge—that is, they become ions. Sodium chloride (NaCl) is the classic ionic material. A sodium atom can assume the electron structure of neon (and a net single positive charge with a reduction in size) by a transfer of its one va- lence 3s electron to a chlorine atom (Figure 2.11a). After such a transfer, the chlorine ion acquires a net negative charge, an electron configuration identical to that of argon; it is also larger than the chlorine atom. Ionic bonding is illustrated schematically in Figure 2.11b. The attractive bonding forces are coulombic—that is, positive and negative ions, by virtue of their net electrical charge, attract one another.

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  eBook - PDF

  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)

  In addition, these free electrons act as a ‘glue to hold the ion cores together. Bonding energies and melting temperatures for several metals are listed in table 2.3. Bonding may be weak or strong; energies range from 62 kJ/mol for mercury to 850 kJ/mol for tungsten. Their respective melting temperatures are −39°C and 3414°C (−39°F and 6177°F). FIGURE 2.19 Schematic illustration of Metallic Bonding. - - - - - - - - - - - - - - - - - - - - - - + + + + + + + + + + + + + + + + + + + + + Ion core Metal atom Sea of valence electrons - - Metallic Bonding is found in the periodic table for Group IA and IIA elements and, in fact, for all elemental metals. Metals are good con- ductors of both elec- tricity and heat as a consequence of their free electrons. Further- more, we have noted that at room tempera- ture, most metals and their alloys fail in a ductile manner — that is, fracture occurs after the materials have experienced significant degrees of permanent deformation. This behaviour is explained in terms of a deformation mechanism, which is implicitly related to the characteristics of the metallic bond. 2.7 Secondary bonding or van der Waals bonding Secondary bonds, or van der Waals (physical) bonds, are weak in comparison to the primary or chemical bonds; bonding energies range between about 4 and 30 kJ/mol. Secondary bonding exists between virtually all atoms or molecules, but its presence may be obscured if any of the three primary bonding types is present. Secondary bonding is evidenced for the inert gases, which have stable electron structures. In 34 Materials science and engineering: an introduction addition, secondary (or intermolecular) bonds are possible between atoms or groups of atoms, which themselves are joined together by primary (or intramolecular) ionic or covalent bonds. Secondary bonding forces arise from atomic or molecular dipoles.

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  Solid State Materials Chemistry

  • Patrick M. Woodward, Pavel Karen, John S. O. Evans, Thomas Vogt(Authors)
  • 2021(Publication Date)
  • Cambridge University Press

    (Publisher)

  5 Chemical Bonding Changes in crystal structure invariably lead to changes in physical and/or chemical proper- ties. In some cases, these changes can be dramatic, as illustrated by the contrasting properties of the allotropes of carbon (diamond, graphite, graphene, C 60 , etc.); in other cases they are subtle but nonetheless important. To understand the relationship between structure and properties, one must first understand chemical bonding. We begin this chapter with an overview of ionic bonding. From there we move on to the properties of atomic orbitals (AOs) and their interactions to form covalent bonds through the framework of molecular orbital theory. In Chapter 6, we then build upon these principles to describe the formation of bands in extended solids. In this way, covalent and Metallic Bonding can be understood through a common approach. 5.1 Ionic Bonding Although there are no compounds where the bonding can be described as purely ionic, the ionic model is a useful approximation for many compounds. We begin our treatment of bonding with a brief overview of the factors that determine the strength of ionic bonding in crystalline solids. 5.1.1 Coulombic Potential Energy The coulombic potential energy, U C , between two ions of charge numbers z 1 and z 2 separated by a distance d is: U C ¼ ðz 1 eÞ  ðz 2 eÞ 4πε 0 d (5.1) 154 where e is the elementary charge and ε 0 is the electric constant. 1 To estimate the strength of ionic bonding in a crystal, we treat the ions as point charges and use Equation (5.1) to capture all electrostatic interactions in the crystal, both attractive and repulsive. To illustrate, consider the electrostatic interactions in the NaCl structure shown in Figure 5.1. We begin with the Cl − ion in the center of the unit cell and consider the interaction between this ion and all other ions in the crystal.

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  The Electronic Structures of Solids

  • B. R. Coles, A. D. Caplin(Authors)
  • 2013(Publication Date)
  • Arnold

    (Publisher)

  2 Bonding Between Atoms 2.1 Introduction When we form an elemental solid by condensation from a gas of weakly interacting atoms the bonding energy arises from modifications in the states of the outer electrons, such that they are accommodated at lower energies. In simple materials that are solid at room temperature these electrons are a small number (per atom) of well-defined electrons which are termed the valency electrons, since they are those that take part in chemical bonding processes with other types of atoms to form molecules; an understanding of such chemical bonding processes is clearly a valuable step towards an understanding of the electronic structures of solids. Unfortunately many students of physics nowadays have only limited awareness of chemistry and we shall not assume any detailed chemical knowledge beyond the following: (i) On the left hand side of the periodic table elements in the first three columns have valencies equal to their group number, and tend to form ionic compounds in which they are positive ions with closed shells. (ii) Elements on the right hand side of the periodic table in groups VI and VII have valencies equal to 8 minus the group number, and tend to form ionic compounds in which they are negative ions with closed shells. (iii) Elements in the transition groups show various valencies; the outer s-electrons are always involved but no simple rules can be given for the involvement of the d-electrons in bonding. (iv) Elements in the rare-earth group are almost always trivalent in compounds with the main exceptions of Eu (often divalent) Yb (sometimes divalent) and Ce (sometimes 4-valent). These features are not unreasonable in the light of the general characteristics of atomic energy levels that we discussed in the previous chapter, and often in the solid there is an important contribution to cohesion from the Coulomb attraction between charged ions of opposite sign.

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