Chemistry
Bond Length
Bond length refers to the distance between the nuclei of two bonded atoms. It is a crucial factor in determining the strength and stability of a chemical bond. The bond length is influenced by the types of atoms involved and the nature of the bond, such as single, double, or triple bonds.
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3 Key excerpts on "Bond Length"
eBook - PDF- Edward J. Neth, Paul Flowers, Klaus Theopold, Richard Langley, William R. Robinson(Authors)
- 2019(Publication Date)
- Openstax(Publisher)
However, molecular structure is actually three-dimensional, and it is important to be able to describe molecular bonds in terms of their distances, angles, and relative arrangements in space ( Figure 4.14). A bond angle is the angle between any two bonds that include a common atom, usually measured in degrees. A bond distance (or Bond Length) is the distance between the nuclei of two bonded atoms along the straight line joining the nuclei. Bond distances are measured in Ångstroms (1 Å = 10 –10 m) or picometers (1 pm = 10 –12 m, 100 pm = 1 Å). FIGURE 4.14 Bond distances (lengths) and angles are shown for the formaldehyde molecule, H 2 CO. VSEPR Theory Valence shell electron-pair repulsion theory (VSEPR theory) enables us to predict the molecular structure, including approximate bond angles around a central atom, of a molecule from an examination of the number of bonds and lone electron pairs in its Lewis structure. The VSEPR model assumes that electron pairs in the valence shell of a central atom will adopt an arrangement that minimizes repulsions between these electron pairs by maximizing the distance between them. The electrons in the valence shell of a central atom form 206 4 • Chemical Bonding and Molecular Geometry Access for free at openstax.org either bonding pairs of electrons, located primarily between bonded atoms, or lone pairs. The electrostatic repulsion of these electrons is reduced when the various regions of high electron density assume positions as far from each other as possible. VSEPR theory predicts the arrangement of electron pairs around each central atom and, usually, the correct arrangement of atoms in a molecule. We should understand, however, that the theory only considers electron- pair repulsions. Other interactions, such as nuclear-nuclear repulsions and nuclear-electron attractions, are also involved in the final arrangement that atoms adopt in a particular molecular structure.
eBook - PDF- Paul Flowers, Klaus Theopold, Richard Langley, William R. Robinson(Authors)
- 2019(Publication Date)
- Openstax(Publisher)
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. 7.2 • Covalent Bonding 317 FIGURE 7.4 The potential energy of two separate hydrogen atoms (right) decreases as they approach each other, and the single electrons on each atom are shared to form a covalent bond. The Bond Length is the internuclear distance at which the lowest potential energy is achieved. It is essential to remember that energy must be added to break chemical bonds (an endothermic process), whereas forming chemical bonds releases energy (an exothermic process). In the case of H 2 , the covalent bond is very strong; a large amount of energy, 436 kJ, must be added to break the bonds in one mole of hydrogen molecules and cause the atoms to separate: Conversely, the same amount of energy is released when one mole of H 2 molecules forms from two moles of H atoms: Pure vs. Polar Covalent Bonds If the atoms that form a covalent bond are identical, as in H 2 , Cl 2 , and other diatomic molecules, then the electrons in the bond must be shared equally. We refer to this as a pure covalent bond. Electrons shared in 318 7 • Chemical Bonding and Molecular Geometry Access for free at openstax.org pure covalent bonds have an equal probability of being near each nucleus. In the case of Cl 2 , each atom starts off with seven valence electrons, and each Cl shares one electron with the other, forming one covalent bond: The total number of electrons around each individual atom consists of six nonbonding electrons and two shared (i.e., bonding) electrons for eight total electrons, matching the number of valence electrons in the noble gas argon.
eBook - PDF- Brian W. Pfennig(Author)
- 2021(Publication Date)
- Wiley(Publisher)
As is often the case, the individual properties of atoms or groups of atoms (such as size, charge, Bond Length, and energy) change depending on the context in which they TABLE 4.2 Average Bond Lengths and bond dissociation energies (BDE) for selected types of covalent bonds. Bond Type Ave Bond Length (pm) Ave BDE (kJ/mol) Bond Type Ave Bond Length (pm) Ave BDE (kJ/mol) H─H 74 432 B─F 130 732 C─H 109 435 N─N 145 167 N─H 101 386 N─O 140 201 O─H 96 459 N─F 136 283 Si─H 148 318 P─P 221 201 P─H 144 322 P─O 163 335 S─H 134 363 P─F 154 490 H─F 92 565 P─Cl 203 326 H─Cl 127 428 O─O 148 142 H─Br 141 362 O─F 142 190 H─I 161 295 S─F 156 284 F─F 142 155 C=C 134 602 Cl─Cl 199 240 C= ─ C 120 835 Br─Br 228 190 C=N 129 615 I─I 267 148 C= ─ N 116 887 C─C 154 346 C=O 120 799 C─N 147 305 C= ─ O 113 1072 C─O 143 358 N=N 125 418 C─F 135 485 N= ─ N 110 942 C─S 182 272 N=O 121 607 C─Cl 177 327 O=O 121 494 C─Br 194 285 P=O 150 544 B─Cl 175 427 S=O 143 522 143 4.4 COVALENT Bond LengthS AND BOND DISSOCIATION ENERGIES are found. According to Table 4.2, for example, a typical C─H bond will require about 411 kJ/mol of energy on average to break. However, the BDE for a C─H bond in CH 4 is 435 kJ/mol, while it is only 380 kJ/mol in Br 3 CH, where the more electronegative Br atoms withdraw more of the molecule’s electron density away from the C─H bond. The exact bond dissociation energy depends on a variety of factors, including the degree of orbital overlap, the diffuseness and polarizability of the overlapping orbitals, the number of electron pairs between the atoms, crystal packing effects, and the percent ionic character. For example, the Bond Lengths in the halogens F 2 , Cl 2 , Br 2 , and I 2 increase down the column as the radius of the individual atoms increases. At the same time, the bond dissociation energies for the halogens decrease down the column as the orbitals become more diffuse (with fluorine being an exception due to the uniqueness principle).
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