6: Chemical Bonding - Electron Pairs and Octets
- Page ID
- 49340
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Theories of chemical bonding invariably involve electrons. When one atom approaches another, the valence electrons, found in the outermost regions of the atoms, interact long before the nuclei can come close together. Electrons are the least massive components of an atom, and so they can relocate to produce electrostatic forces which hold atoms together. According to Coulomb’s law, such electrostatic or Coulombic forces are quite large when charges are separated by distances of a few hundred picometers—the size of an atom. Coulombic forces, then, are quite capable of explaining the strengths of the bonds by which atoms are held together.
- 6.1: Prelude to Chemical Bonding
- Theories of chemical bonding invariably involve electrons. When one atom approaches another, the valence electrons, found in the outermost regions of the atoms, interact long before the nuclei can come close together. Electrons are the least massive components of an atom, and so they can relocate to produce electrostatic forces which hold atoms together.
- 6.2: Ionic Bonding
- Ionic bonding involves transfer of an electron from one atom (which becomes a positively charged cation) to another (which becomes a negatively charged anion). The two ions attract strongly to form a crystal lattice.
- 6.3: Energy and the Formation of Ions
- Formation of an ion pair by transfer of an electron from an Li atom to an H atom results in an overall lowering of the total energy of the two nuclei and four electrons involved.
- 6.4: The Ionic Crystal Lattice
- The formation of such an ionic crystal lattice results in a lower potential energy than is possible if the ions only group into pairs.
- 6.5: Ions and Noble-Gas Electron Configurations
- Ions often form in characteristic ways, aiming to achieve noble gas configuration.
- 6.6: Ionization Energies
- If the ionization energies of the elements are plotted against atomic number, an obvious feature is observed whereby elements with the highest ionization energies are the noble gases. Since the ionization energy measures the energy which must be supplied to remove an electron, these high values mean that it is difficult to remove an electron from an atom of a noble gas.
- 6.7: Ionization of Transition and Inner Transition Elements
- Furthermore, experimental measurements show that for transition and inner transition elements the electrons lost when ionization occurs are not the last ones which were added to build up the atomic electron configuration. Instead, electrons are usually removed first from the subshell having the largest principal quantum number.
- 6.8: Electron Affinities
- Electron affinities are more difficult to measure experimentally than are ionization energies, and far fewer values are available. The relationship of the periodic table with those electron affinities that have been measured or estimated from calculations can be seen on the table of ionization energies and electron affinities, seen below.
- 6.9: Binary Ionic Compounds and Their Properties
- All ionic compounds have numerous properties in common. Consequently, the ability to recognize an ionic compound from its formula will allow you to predict many of its properties. This is often possible in the case of a binary compound (one which contains only two elements), because formation of a binary ionic compound places quite severe restrictions on the elements involved.
- 6.10: The Octet Rule
- A convenient method for doing this is to regard the compound as being formed from its atoms and to use Lewis diagrams. The octet rule can then be applied. Each atom must lose or gain electrons in order to achieve an octet. Furthermore, all electrons lost by one kind of atom must be gained by the other.
- 6.11: Physical Properties
- Ionic compounds have certain physical characteristics that distinguish them from other compounds. Physical factors like melting point, solubility, and more are discussed in relation to ionic compounds.
- 6.12: Chemical Properties
- The most important chemical characteristic of ionic compounds is that each ion has its own properties. Such properties are different from those of the atom from which the ion was derived.
- 6.13: The Covalent Bond
- Formation of an ionic bond by complete transfer of an electron from one atom to another is possible only for a fairly restricted set of elements. Covalent bonding, in which neither atom loses complete control over its valence electrons, is much more common. In a covalent bond the electrons occupy a region of space between the two nuclei and are said to be shared by them.
- 6.14: Covalent Molecules and the Octet Rule
- The idea that a molecule could be held together by a shared pair of electrons was first suggested by Lewis in 1916. Although Lewis never won the Nobel prize for this or his many other theories, the shared pair of electrons is nevertheless one of the most significant contributions to chemistry of all time. Wave mechanics was still 10 years in the future, and so Lewis was unable to give any mathematical description of exactly how electron sharing was possible.
- 6.15: Writing Lewis Structures for Molecules
- Lewis structures, while rudimentary, allow scientists to quickly display a compound and make inferences as to its 3D structure.
- 6.16: Examples of Lewis Structures
- Lewis structures can be tricky, and this page contains a variety of examples to improve your understanding of how they work using a variety of different elements.
- 6.17: Polyatomic Ions
- Polyatomic ions, common in any lab, contain several atoms covalently bonded together. Often, these ions are charged and combine with metals to form ionic bonds.
- 6.18: Ionic Compounds Containing Polyatomic Ions
- Polyatomic ions are everywhere and this pages introduces you to familiar polyatomic ions that often form ionic bonds.
- 6.19: Atomic Sizes
- Atomic size changes in predictable ways as one moves around the periodic table. In this section, we learn the periodic trends of atomic size.
- 6.20: Ionic Sizes
- The loss or addition of electrons (aka ionization) changes atomic size. Read on to find out how and why.