Valence electrons in molecules

Lewis structures are blueprints that show the distribution of valence electrons in molecules. However, the dots and lines of a Lewis structure do not show any details of how bonds form, how molecules react, or the shape of a molecule. In this respect, a Lewis structure is like the electron configuration of an atom both tell us about electron distributions, but neither provides detailed descriptions. Just as we need atomic orbitals to understand how electrons are distributed in an atom, we need an orbital view to understand how electrons are distributed in a molecule. 

In each of our examples, only bonding pairs of electrons surrounded the central atom. Many times the central atom contains a single lone pair or lone pairs of electrons in addition to bonding pairs. The presence of lone pairs, which occupy more space than bonding pairs, affects the repulsive forces between the valence electrons in molecules. 

Thus, A is the correction for pairs of valence electrons in molecules, B the correction for unpaired electrons in molecules, C the correction for pairs of valence electrons in atoms, and D the correction for unpaired electrons in atoms. The use of different corrections for atoms and molecules can be justihed, in part, by noting that effects of basis functions with higher angular momentum are likely to be of more importance in molecules than in atoms. The A, B, C, D values are chosen to give the smallest average absolute deviation from experiment for the G2/97 test set. For G3 theory, A = 6.386 mhartrees, B = 2.977 mhartrees, C = 6.219 mhartrees, D = 1.185 mhartrees. 

Lewis presented a simple, but useful, method of describing the arrangement of valence electrons in molecules. The approach uses dots (or dots and crosses) to represent the number of valence electrons, and the nuclei are indicated by appropriate elemental S5anbols. A basic premise of the theory is that electrons in a molecule should be paired the presence of a single (odd) electron indicates that the species is a radical. 

The functionals considered in Section 1.3 are all semilocal the local kinetic energy at the point r depends only on the electron density and its derivatives at the point r. Improved models for the kinetic energy require considering how the electron density at other points r affects the local kinetic energy at the point r. Without including these effects, the oscillations in electron density that are essential for modeling the shell structure of atoms and differentiating between core and valence electrons in molecules cannot be recovered. 

In Sections 7-1 and 7-2 we drew Lewis dot formulas for atoms and monatomic ions. We can use Lewis formulas to show the valence electrons in molecules and polyatomic ions. 

Electron-dot formulas can be used to diagram the sharing of valence electrons in molecules and polyatomic ions. The presence of multiple bonds can be identified, and possible resonance structures can be drawn. From the electron-dot formulas, we can predict the three-dimensional shapes and polarities of molecules. Then we examine how the different attractive forces between the particles of ions and molecules influence their physical properties, such as melting and boiling point. Finally, we discuss the physical states of solids, liquids, and gases and describe the energy involved in changes of state. 

The valence electrons in molecules are shown using an electron-dot formula, also called a Lewis structure. The shared electrons, or bonding pairs, are shown as two dots or a single line between atoms. The nonbonding pairs of electrons, or lone pairs, are placed on the outside. For example, a fluorine molecule, F2, consists of two fluorine atoms, which are in Group 7A (17), each with seven valence... 

The term resonance has also been applied in valency. The general idea of resonance in this sense is that if the valency electrons in a molecule are capable of several alternative arrangements which differ by only a small amount in energy and have no geometrical differences, then the actual arrangement will be a hybrid of these various alternatives. See mesomerism. The stabilization of such a system over the non-resonating forms is the resonance energy. 

The sum over all entries of the BE-matrix (S ) gives the total number of valence electrons in the molecule (Eq. (2)). 

The simplest many-electron wave function that satisfies the Exclusion Principle is a product of N different one-electron functions that have been antisymmetrized, or written as a determinant. Here, N is the number of electrons (or valence electrons) in the molecule. HyperChem uses this form of the wave function for most semi-empirical and ab initio calculations. Exceptions involve using the Configuration Interaction option (see page 119). HyperChem computes one-electron functions, termed molecular spin orbitals, by relatively simple integration and summation calculations. The many-electron wave function, which has N terms (the number of terms in the determinant), never needs to be evaluated. 

UV/Vis Spectra for Molecules and Ions When a molecule or ion absorbs ultraviolet or visible radiation it undergoes a change in its valence electron configuration. The valence electrons in organic molecules, and inorganic anions such as oc-... 

To express the calculations in a general way, the formal charge on an atom is equal to the number of valence electrons in a neutral, isolated atom minus the number of electrons owned by that atom in a molecule. The number of electrons in the bonded atom, in turn, is equal to half the number of bonding electrons plus the nonbonding, lone-pair electrons. 

This idea is readily extended to simple molecules of compounds formed by nonmetal atoms. An example is the HF molecule. You will recall that a fluorine atom has the electron configuration ls22s22p5. ft has seven electrons in its outermost principal energy level (n = 2). These are referred to as valence electrons, in contrast to the core electrons filling the principal level, n = 1. If the valence electrons are shown as dots around the symbol of the element, the fluorine atom can be represented as... 

Notice that in each case the oxygen or nitrogen atom is surrounded by eight valence electrons. In each species, a single electron pair is shared between two bonded atoms. These bonds are called single bonds. There is one single bond in the OH- ion, two in the H20 molecule, three in NH3, and four in NH4+. There are three unshared pairs in the hydroxide ion, two in the water molecule, one in the ammonia molecule, and none in the ammonium ion. 

The largest class of molecules to violate the octet rule consists of species in which the central atom is surrounded by more than four pairs of valence electrons. Typical molecules of this type are phosphorus pentachloride, PC15, and sulfur hexafluoride, SF6. The Lewis structures of these molecules are... 

The valence electrons in a molecule are distributed among the available molecular orbitals. The process followed is much like that used with electrons in atoms. In particular, we find the following ... 

Another species in which delocalized pi orbitals play an important role is benzene, QHg. There are 30 valence electrons in the molecule, 24 of which are required to form the sigma H H bond framework ... 

There is another possible consequence of a collision between two fluorine atoms. The two atoms can remain together to form a molecule. Each atom has a valence electron in a half-filled orbital. We can imagine these two atoms orienting so that these half-filled" orbitals overlap in space. Then the half-filled" valence orbital of... 

In the molecular orbital description of homonuclear diatomic molecules, we first build all possible molecular orbitals from the available valence-shell atomic orbitals. Then we accommodate the valence electrons in molecular orbitals by using the same procedure we used in the building-up principle for atoms (Section 1.13). That is,... 

When N valence atomic orbitals overlap, they form N molecular orbitals. The ground-state electron configuration of a molecule is deduced by using the building-up principle to accommodate all the valence electrons in the available molecular orbitals. The bond order is the net number of bonds that hold the molecule together. 

Representations showing electrons in molecules seem to suggest localisation of the valence electrons, but there are problematic issues in this regard. For example, we might ask if dioxygen has a double bond and two lone pairs on each O atom (as in Table 1.1) - a stmcture that does not reconcile with the paramagnetic nature of the substance - or a single bond and an odd number of electrons localised on each atom, as shown here ... 

In this section, we develop a process for making schematic drawings of molecules called Lewis structures. A Lewis structure shows how the atoms in a molecule are bonded together. A Lewis structure also reveals the distribution of bonding and nonbonding valence electrons in a molecule. In a sense, a Lewis structure is a molecular blueprint that... 

A Lewis base must have valence electrons available for bond formation. Any molecule whose Lewis stmcture shows nonbonding electrons can act as a Lewis base. Ammonia, phosphorus trichloride, and dimethyl ether, each of which contains lone pairs, are Lewis bases. Anions can also act as Lewis bases. In the first example of adduct formation above, the fluoride ion, with eight valence electrons in its 2 s and 2 p orbitals, acts as a Lewis base. 

A soft Lewis base has a large donor atom of high polarizability and low electronegativity. Iodide ion has its valence electrons in large a = 5 orbitals, making this anion highly polarizable and a very soft base. Other molecules and polyatomic anions with donor atoms from rows 3 to 6 are also soft bases. To summarize, the donor atom becomes softer from top to bottom of a column of the periodic table. 

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