Chemistry
Bonding and Elemental Properties
Bonding in chemistry refers to the attractive forces that hold atoms together in compounds. These bonds can be covalent, where atoms share electrons, or ionic, where electrons are transferred from one atom to another. Elemental properties are the characteristics of individual elements, such as their reactivity, melting point, and conductivity, which are determined by their atomic structure and bonding behavior.
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11 Key excerpts on "Bonding and Elemental Properties"
- 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.
eBook - PDF- 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.
eBook - PDFFundamentals 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.
eBook - PDF- Bernard Moody(Author)
- 2013(Publication Date)
- Arnold(Publisher)
5 Bonding and the structures displayed by elements and their compounds The general physical properties of compounds related to bond type The nature of the bonding in a compound and of the geometrical pattern adopted by the ions or molecules in a solid, will largely determine the physical properties ofthat substance. While distinc-tive properties associated with ionic and covalent bonding may be discerned, there is a gradual merging of characteristics when the compounds of a large number of elements are compared. This gradual transition is not altogether unexpected. When fused or dissolved in water, an ionic com-pound will conduct electricity. The current is carried through the liquid by the ions which gain their mobility when the compound is melted or dispersed in a solvent. Unless a reaction occurs with the solvent, covalent substances yield non-conducting liquids. In the crystal lattice of an ionic compound, each ion is surrounded by oppositely charged ions, the number depending on the particular pattern adopted in the crystal. Strong electrical forces hold the ions in position although each atom oscillates by virtue of its thermal energy. Considerable energy is required to overcome the forces of attraction and ionic compounds usually melt at high temperatures and are non-volatile. On the other hand, covalent molecules, each electrically neutral, are held by much weaker intermolecular forces. Therefore, fusion, boiling and sublimation are relatively easy to accomplish. To illustrate this point, the melting-points of the fluorides formed by the elements of Period 3, sodium-sulphur, are shown in Table 5.1.
eBook - PDF- 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.
Alternatively, atoms may be pushed into new configurations by adding energy to systems. Much of industrial chemis- try, from the smelting of iron to the synthesis of plastics, operates on this principle. 1 H 1.00794 3 Li 6.941 4 Be 9.01218 11 Na 22.98977 12 Mg 24.3050 2 He 4.00260 5 B 10.811 6 C 12.011 7 N 14.00674 8 O 15.9994 9 F 18.99840 10 Ne 20.1797 13 Al 26.98154 14 Si 28.0855 15 P 30.97376 16 S 32.066 17 Cl 35.4527 18 Ar 39.948 Figure 10-1 • The first three rows of the periodic table, containing elements 1 and 2, 3 through 10, and 11 through 18, respectively, hold the key to understanding chemical bonding. Types of Chemical Bonds Atoms link together by three principal kinds of chemical bonds—ionic, metallic, and covalent—all of which involve redistributing electrons between atoms. In addition, polar- ization, hydrogen bonding, and van der Waals forces result from shifts of electrons within their atoms or groups of atoms. Each type of bonding corresponds to a different way of rearranging electrons, and each produces distinctive properties in the materials it forms. Types of Chemical Bonds | 209 IONIC BONDS We’ve seen that atoms with “magic numbers” of 2, 10, 18, or 36 electrons are particu- larly stable. By the same token, atoms that differ from these magic numbers by only one electron in their outer orbits are particularly reactive—in effect, they are “anxious” to fill or empty their outer orbits. Such atoms tend to form ionic bonds, chemical bonds in which the electrical force between two oppositely charged ions holds the atoms together. Ionic bonds often form as one atom gives up an electron while another receives it. Sodium (a soft, silvery white metal), for example, has 11 electrons in an electrically neutral atom—2 in the lowest orbit, 8 in the next, and a single electron with lots of chemical potential energy in its outer shell. Sodium’s best bonding strategy, therefore, is to lose one electron.
eBook - PDFFundamentals of Materials Science and Engineering
An Integrated Approach
- William D. Callister, Jr., David G. Rethwisch(Authors)
- 2016(Publication Date)
- Wiley(Publisher)
(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. Each atom consists of a very small nucleus composed of protons and neutrons and is encircled by moving electrons. 1 Both electrons and protons are electrically charged, the charge magnitude being 1.602 × 10 −19 C, which is negative in sign for electrons and posi- tive for protons; neutrons are electrically neutral. Masses for these subatomic particles are extremely small; protons and neutrons have approximately the same mass, 1.67 × 10 −27 kg, which is significantly larger than that of an electron, 9.11 × 10 −31 kg. Each chemical element is characterized by the number of protons in the nucleus, or the atomic number (Z). 2 For an electrically neutral or complete atom, the atomic number also equals the number of electrons. This atomic number ranges in integral units from 1 for hydrogen to 92 for uranium, the highest of the naturally occurring elements. The atomic mass (A) of a specific atom may be expressed as the sum of the masses of protons and neutrons within the nucleus.
eBook - PDFFundamentals of Materials Science and Engineering
An Integrated Approach
- William D. Callister, Jr., David G. Rethwisch(Authors)
- 2022(Publication Date)
- Wiley(Publisher)
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. bonding energy primary bond 2.6 | | PRIMARY INTERATOMIC BONDS 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 2.13a). After such a transfer, the chlorine ion acquires a net negative charge, an electron configuration identical to that of argon; ionic bonding 2.6 Primary Interatomic Bonds 35 it is also larger than the chlorine atom. Ionic bonding is illustrated schematically in Figure 2.13b. The attractive bonding forces are coulombic—that is, positive and negative ions, by virtue of their net electrical charge, attract one another. For two isolated ions, the attrac- tive energy E A is a function of the interatomic distance according to E A = − A __ r (2.9) Theoretically, the constant A is equal to A = 1 ____ 4 π ε 0 (|Z 1 |e)(|Z 2 |e) (2.10) Here ε 0 is the permittivity of a vacuum (8.85 × 10 −12 F/m), |Z 1 | and |Z 2 | are absolute values of the valences for the two ion types, and e is the electronic charge (1.602 × 10 −19 C).
eBook - ePubUnderstanding Basic Chemistry
The Learner's Approach
- Kim Seng Chan, Jeanne Tan(Authors)
- 2014(Publication Date)
- WSPC(Publisher)
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 of valence electrons available for bonding. (This factor is useful for explaining the increase in metallic bond strength across a period of metals. For example, from Na to Mg to Al.)•The size of the metal cation. (This factor is useful for explaining the decrease in metallic bond strength down a group of metals. For example, from Na to K to Rb to CS.)3.1.1 Physical properties of metalsThe physical properties of a substance are related to the structure and the types of bonding present in the structure. There are three important aspects to focus on:•Volatility The volatility of a compound is determined by the strength of the bond between the particles. The strength of the bond can be measured by measuring the melting point/boiling point of the compound. The higher the melting point, the stronger the bond between the particles. •Conductivity The conductivity of a compound is determined by whether the compound contains mobile charge carriers. A charge carrier can be ions or electrons. Thus, conductivity is a measurement of whether the compound conducts electricity readily or not. •Solubility The nature of the solvent is important when we discuss solubility. A solvent can be polar or non-polar. Later, we would discuss more about how a molecule can be polar under Section 3.3 on Covalent Bonding. Volatility of metalsGenerally, a high melting point/boiling point is due to the strong electrostatic attractive force between the positive ions and delocalized valence electrons. Thus, a lot of heat energy is required to break these forces.For instance, the higher melting point of Al (m.p. 660°C) as compared to Mg (m.p. 650°C) is due to greater number
eBook - PDFMaterials 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
eBook - PDF- Edward J. Neth, Paul Flowers, Klaus Theopold, Richard Langley, William R. Robinson(Authors)
- 2019(Publication Date)
- Openstax(Publisher)
For example, two hydrogen atoms bond covalently to form an H 2 molecule; each hydrogen atom in the H 2 molecule has two electrons stabilizing it, giving each atom the same number of valence electrons as the noble gas He. Compounds that contain covalent bonds exhibit different physical properties than ionic compounds. Because the attraction between molecules, which are electrically neutral, is weaker than that between electrically charged ions, covalent compounds generally have much lower melting and boiling points than ionic compounds. In fact, many covalent compounds are liquids or gases at room temperature, and, in their solid states, they are typically much softer than ionic solids. Furthermore, whereas ionic compounds are good conductors of electricity when dissolved in water, most covalent compounds are insoluble in water; since they are electrically neutral, they are poor conductors of electricity in any state. Formation of Covalent Bonds Nonmetal atoms frequently form covalent bonds with other nonmetal atoms. For example, the hydrogen molecule, H 2 , contains a covalent bond between its two hydrogen atoms. Figure 4.4 illustrates why this bond is formed. Starting on the far right, we have two separate hydrogen atoms with a particular potential energy, indicated by the red line. Along the x-axis is the distance between the two atoms. As the two atoms approach each other (moving left along the x-axis), their valence orbitals (1s) begin to overlap. The single electrons on each hydrogen atom then interact with both atomic nuclei, occupying the space around both atoms. The strong attraction of each shared electron to both nuclei stabilizes the system, and the potential energy decreases as the bond distance decreases. If the atoms continue to approach each other, the positive charges in the two nuclei begin to repel each other, and the potential energy increases. The bond length is determined by the distance at which the lowest potential energy is achieved.
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