AMERICAN JOURNAL OF UNDERGRADUATE RESEARCH VOL. 5, NO. 1 (2006) Visualizing Bond Types with Electron Density Models: How Informative is Electronegativity? Heath D. Stotts and J. Conceicao* Division of Natural Sciences Northwestern Oklahoma State University 709 Oklahoma Blvd. Alva, Oklahoma 73717 USA Received: March 2, 2006 Accepted: April 4, 2006 ABSTRACT Electron densities are used to visualize pure covalent, polar covalent and ionic bonds in binary compounds. The rationale for this study stems in part from the observations that within the same bond type, for example pure covalent, a variety of bond properties exist. Simple ∆EN predictions by Pauling do not adequately explain differences within the same bond type, nor determine covalent or ionic bonding. In this study, a series of electron density maps for binary compounds have been calculated to compare the characteristics of the maps to ∆EN predictions. I. INTRODUCTION The modern description of bond types is based on the relative The concept of chemical bonding is electronegativity of the atoms involved in the by far one of the most useful, and at the bond. Bond type can be predicted using same time one of the most difficult to different models; the most prevalent one has understand in all of chemistry [1]. been Pauling’s difference of electro- Understanding the structure and connectivity negativity (∆EN) and is found to be accurate in bonds may allow us to better predict the in most cases [3]. An exception to Pauling’s properties of compounds and lead to the ∆EN is the LiI molecule which is predicted to development of new materials, polymers, be polar covalent, but experimentally it is and other advanced technologies such as ionic in character. An alternative method thin films, synthetic fibers and artificial blood. has been proposed by Sproul [4]. According The importance of this concept was to his model, a more complete description of recognized by van Arkel who classified a bond type requires the use of an additional bond as one of three types, these being: parameter, this being the average of the pure covalent, polar covalent, and ionic [2]. electronegativities, ½ (χA + χB),B rather than Specifically, a pure covalent bond exhibits simply ∆EN. Both models though are still equally shared electron density between the incomplete; for example in the pure covalent two bonding atoms with no net dipole bonds of H2 or Cl2, differences still exist in moment. A polar covalent bond exhibits a bond energies. Bond typing based solely on shift in the bonding electron density and has electronegativity (∆EN and ½ (χA + χB))B is a net dipole moment. At the furthermost thus inadequate. Ideally, bond type should extreme of the bonding spectrum ionic be determined from electron density bonds are characterized by a lack of measurements. Over the past decade, bonding electron density between the two advancements in experimental and bonding nuclei. In this case, the largest theoretical methods allow accurate dipole moment is expected. determination of electron densities [5-6] for small molecules and crystals [7-8]. With the * The authors may be contacted through their e-mail advent of inexpensive computers and addresses at [email protected] and computational software, electron density [email protected]. 11 AMERICAN JOURNAL OF UNDERGRADUATE RESEARCH VOL. 5, NO. 1 (2006) maps [9-11] are now used routinely and thus reducing the attraction. This diminishes extensively to visualize orbitals, electrostatic electron densities along the bond axis potential, and atomic and molecular size leading to longer and weaker bonds [6]. Our [12]. In this paper we examine the bonding electron density maps correlate to this trend in various binary compounds using electron as is evident in Figure 1. Table 1 is also density models; showing that such maps provided to compare the experimental data may be better in determining bond type of the molecules within the same ∆EN. when compared to relative electronegativity maps like that of Pauling’s ∆EN. b. Polar Covalent II. METHODS AND PROCEDURES Electron density maps for two polar covalent molecules, FI and LiI, are shown in The electron density maps of H2, Cl2, Figure 2. FI and LiI have the same ∆EN of Br2, FI, and LiI respectively exhibit 1.5 and are classified simply as polar isosurfaces at 0.08 au [6] and calculated at covalent bonds according to Pauling’s the Hartree-Fock level using the 3-21G* predictions; the electron densities of these basis set in the Spartan ’04 program. 2-D two compounds tell a very different story electron density contour maps for KCl and though in regards to their respective bonding RbCl were calculated at the same level of natures. In LiI, the lack of electron density in theory using the same basis sets in the the bonding region is representative of an HyperChem 7.5 Student’s Edition [13]. ionic bonding situation, while the FI Geometry optimizations were performed on compound shows an asymmetric sharing of all the compounds prior to the calculations of electrons between the bonding nuclei electron density. indicative of a polar covalent bond. Furthermore, the large differences in dipole III. RESULTS AND DISCUSSION moments and bond energy (Table 2) also support the electron density predictions of a. Pure Covalent Figure 2. Indeed, concluding FI [4] as polar covalent compounds and LiI [11] as ionic. The calculated electron density To understand this correlation, one maps for a series of molecules (H2, Cl2 and can invoke simple bonding ideas involving Br2) with pure covalent bonds are shown in orbital interaction. The energy of the valence Figure 1. All of the molecules display a 2p (≈ -15 hartrees) [15] orbital of fluorine symmetric distribution of electron density and the valence 5p (≈ -20 hartrees) orbital of between the bonding atoms. For H2 the iodine are comparable in magnitude, and greatest electron density is localized thus are expected to interact favorably to between the bonding atoms while the form a covalent bond in FI. Since the atomic diatomic species in the halogen group show orbital of fluorine is lower in energy, the a distinct dip/waist, or lower quantity of bonding molecular orbital containing the electrons between the bonding atoms. electron pair would be fluorine-like, resulting A pure covalent bond is one that is in a slight negative charge about the fluorine stated as having ∆EN of 0; although this is nucleus. LiI on the other hand is quite correct there are some discrepancies in this different. The energy of the 2s (≈ -5 broad statement when referring to the hartrees) valence orbital of lithium and the different experimental bond energies of each 5p (≈ -20 hartrees) orbital of iodine differ molecule (Table 1). Our study shows that substantially and thus covalent bond the bonding electron densities for H2, Cl2 formation between these orbitals diminishes. and Br2 vary with increasing atom radii from The bonding molecular orbital would be H to Br reflecting different degrees of pure more iodine-like and the electron pair would covalency. The lack of a dip/waist in the reside in a region much closer to iodine [16- bonding region of H2, atomic radii of 37 pm 17]. In the case of LiI, the energy mismatch [14], in contrast to a well-defined waist in Br- is large enough that the electron pair resides Br, atomic radii of 114 pm [7], is a exclusively on iodine thereby resulting in the consequence of the increased atomic radii formation of an ionic bond. This justification of the latter. This is expected to place the is consistent with the observations in Figure bonding electrons further from the nuclei 2. In Table 2 experimental data is also 12 AMERICAN JOURNAL OF UNDERGRADUATE RESEARCH VOL. 5, NO. 1 (2006) provided to better compare the characteristics of the two polar covalent molecules. c. Ionic Binary Compounds H 2 The 2-D contour electron density maps of two ionic compounds, KCl and RbCl, are shown in Figure 3. In both instances, polarization of the electron clouds is observed. Both of these binary compounds have a ∆EN of 2.2 and thus are ionic in nature. This simple classification does not explain the discrepancies in the experimental bond energies of 427 KJ/mol for KCl and 448 KJ/mol for RbCl as one Cl2 would expect them to be identical given the apparent assumption that these bonds are the same. Furthermore bond lengths cannot be used to clarify this difference as it depends on the size of the ions. To better understand this observation one has to consider the formation of covalent bonding in ionic compounds (covalency). The inter-nuclear electron cloud in KCl exhibits greater distortion representing a greater degree of Br2 covalent bonding while distortion to a lesser degree is observed in RbCl indicative of Figure 1. Pure Covalent: All models herein smaller covalent character. A greater degree are scaled uniformly to represent their true of covalent character weakens the ionic relative sizes and were computed using the bond in KCl thereby manifesting lower bond 3-21 G* basis set at the Hartree-Fock level energy. In the opposition, RbCl has a lesser and at an isosurface electron density of 0.08 degree of covalent character resulting in au. All models are represented by the same higher bond energy. ∆EN = 0. Dipole Bond Atomic Bond Density Model Molecule ∆EN Moment Energy Radii Length (Debye) (kJ/mol) (pm) H2 0 0 436 37 74 Cl2 0 0 243 99 199 Br2 0 0 193 114 288 Table 1. Comparison of experimental data to density models for H2, Cl2, and Br2. 13 AMERICAN JOURNAL OF UNDERGRADUATE RESEARCH VOL. 5, NO. 1 (2006) an expanded view of the inter-nuclear region.