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General chemistry students need to understand the concept of intermolecular forces and use them to predict the properties of materials. To do this, they must create symbolic representations of mole...

Due to methodological difficulties and limitations of applicability, a quantitative bonding analysis based on the theory of resonance is presently not as convenient and popular as that based on the molecular orbital (MO) methods. Here, we propose an efficient quantitative resonance theory by expanding the DFT wave function in terms of a complete set of Lewis structures. By rigorously separating the resonance subsystem represented by a set of localized MOs, this approach is able to treat large molecules, nonplanar π-conjugate systems, and bonding systems mixing both σ and π electrons. Assessment in 2c-2e systems suggests a new projection-weighted symmetric orthogonalization method to evaluate the weights of resonance contributors, which overcomes the drawbacks of other weighting schemes. Applications to benzene, naphthalene and chlorobenzene show that the present method is insensitive to the basis set employed in the DFT calculations, and to the choices of the independent Lewis set determined by Rumer's rule. Advanced applications to diverse chemical problems provide unique and valuable insights into the understanding of hydrogen bonding, the π substituent effect on benzene, and the mechanism of Diels-Alder reactions.

Latimer and Rodebush (J Am Chem Soc 42: 1419-1433, 1920) discussed the ways a Lewis dot structure could be drawn for liquid water and proposed that the H held between two octets constitutes a bond in 1920. When it was realized that the other molecule of life, DNA, owes its double helix structure to specific hydrogen bonds between A-T (two) and C-G (three) base pairs, the interest in hydrogen bonding grew dramatically. While hydrogen bonding could be readily seen in water and DNA, it was not so easy to understand leading to continuous debates about what it means. This article gives a personal perspective of the evolution of hydrogen bonding since the Latimer and Rodebush paper to the recent IUPAC definition of hydrogen bond, published in 2011 and now. Is there a third C-H center dot center dot center dot O hydrogen bond in the A-T base pair?

Dative Bond (IUPAC): "The coordination bond formed upon interaction between molecular species, one of which serves as a donor and the other as an acceptor of the electron pair to be shared in the complex formed… The distinctive feature of dative bonds is that their minimum-energy rupture in the gas phase or in inert solvent follows the heterolytic bond cleavage path." This definition encompasses an immense number of molecules such as Lewis adducts, transition-metal complexes, supposedly hypervalent or hypovalent systems, and many molecules with multifaceted Lewis structures. Still, there is a large reticence to include dative bonds in the regular depiction of molecules, and even a larger unawareness of the dative bond arrow in many chemical circles. Herein we will discuss in simple chemical terms the past, present and future of such bonds. In addition, we will try to provide cleaner options to represent intricate molecules without sacrificing physical accuracy.

Oxidation number (ON) is taught as an electron-counting concept for redox reactions in chemistry curriculum. The molecular formula method, the Lewis formula method, and the structural formula method have all been used to determine ON. However, the task of assigning ON still poses problems for some teachers and students. This paper explores a new method, the fragmentation method, which is a visual approach for counting the individual ON of any atom according to its structural formula. The critical step is to break the carbon-heteroatom bond into organic fragments and inorganic fragments. The individual ON of carbon atoms and heteroatoms can be determined by the bond cleavages in the organic and biological compounds. The mean ON of carbons can be calculated by the arithmetic mean of all individual ON of carbons in a molecule or molecular ion. The step-by-step operating procedures and examples are provided. When comparing corresponding molecules in organic conversions, the change of individual ON of atoms can be used as a tool for determining the number of transferred electrons. Furthermore, a reaction site can be identified by their changes of individual ON, chemical composition, and bond order in metabolic redox reactions.

Many compounds can be represented well in terms of the two-center–two-electron (2c–2e) bond model. However, it is well-known that this approach has limitations; for example, certain compounds require the use of three-center–two-electron (3c–2e) bonds to provide an adequate description of the bonding. Although a classic example of a compound that features a 3c–2e bond is provided by diborane, B2H6, 3c–2e interactions also feature prominently in transition metal chemistry, as exemplified by bridging hydride compounds, agostic compounds, dihydrogen complexes, and hydrocarbon and silane σ-complexes. In addition to being able to identify the different types of bonds (2c–2e and 3c–2e) present in a molecule, it is essential to be able to utilize these models to evaluate the chemical reasonableness of a molecule by applying the octet and 18-electron rules; however, to do so requires determination of the electron counts of atoms in molecules. Although this is easily achieved for molecules that possess only 2c–2e bonds, the situation is more complex for those that possess 3c–2e bonds. Therefore, this article describes a convenient approach for representing 3c–2e interactions in a manner that facilitates the electron counting procedure for such compounds. In particular, specific attention is devoted to the use of the half-arrow formalism to represent 3c–2e interactions in compounds with bridging hydrogen atoms.

The shared electron-pair bonding model was suggested by Gilbert Lewis more than 100 years ago. Emerging from the chemical experience of the time, Lewis structures described contemporary aspects of chemical reality in terms of empirically adapted models without any (then unknown) quantum physical underpinnings. This Perspective details the origins and historical development of the Lewis model, which we contrast with the physical understanding of chemical bonding in terms of contemporary quantum chemistry. Some intuitively plausible classical explanations of the past, not least of which are the sharing of electrons by two atoms and the subtypes of shared electron-pair bonding and dative bonding, turned out to be well founded. Some other chemical dogmata, including the concept that bonding occurs only between two nuclei and is caused by spin coupling or that bond energy is of purely electrostatic origin, are less well founded. We now know that covalent bonding is not driven by the formation of an electron pair but rather by the lowering of the kinetic energy density of the shared electrons in the bonding region, which is provided by the interference of the atomic wavefunctions. Lewis structures remain highly useful models for describing chemical bonding in molecular structures and chemical reactions, particularly when supported by quantum chemistry. The concepts behind the three most common quantum chemical approximations — the valence bond, molecular orbital and density functional theories — are described. These methods allow us to learn that bonding is an energetic phenomenon, from which descriptors such as bond length, bond dissociation energies and force constants are derivable. The energetic origins of bonding point to bond energy decomposition analysis as a natural tool for elucidating the actions of bonding electrons. Lewis’ shared electron-pair model was a stroke of genius, describing the structure and reactivity of molecules purely on the basis of his tremendous knowledge of empirical chemistry without any quantum chemistry. Unprecedented in simplicity, its success unfortunately concealed some misleading interpretations of the physical origin of chemical bonding.

A partitioning of the wavefunction into tiles allows electronic excitations to be viewed as electron vibrations., Since the conception of the electron pair bond, Lewis structures have been used to illustrate the electronic structure of a molecule in its ground state. But, for excited states, most descriptions rely on the concept of molecular orbitals. In this work we demonstrate a simple and intuitive description of electronic resonances in terms of localized electron vibrations. By partitioning the 3N-dimensional space of a many-electron wavefunction into hyper-regions related by permutation symmetry, chemical structures naturally result which correspond closely to Lewis structures, with identifiable single and double bonds, and lone pairs. Here we demonstrate how this picture of electronic structure develops upon the admixture of electronic wavefunctions, in the spirit of coherent electronic transitions. We show that π–π* transitions correspond to double-bonding electrons oscillating along the bond axis, and n–π* transitions reveal lone-pairs vibrating out of plane. In butadiene and hexatriene, the double-bond oscillations combine with in- and out-of-phase combinations, revealing the correspondence between electronic transitions and molecular normal mode vibrations. This analysis allows electronic excitations to be described by building upon ground state electronic structures, without the need for molecular orbitals.

International audience; This article dwells on the nature of “inverted bonds”, which refer to the σ interaction between two sp hybrids by their smaller lobes, and their presence in [1.1.1]propellane. Firstly, we study H3C−C models of C−C bonds with frozen H‐C‐C angles reproducing the constraints of various degrees of “inversion”. Secondly, the molecular orbital (MO) properties of [1.1.1]propellane and [1.1.1]bicyclopentane are analyzed with the help of orbital forces as a criterion of bonding/antibonding character and as a basis to evaluate bond energies. Triplet and cationic states of [1.1.1]propellane species are also considered to confirm the bonding/antibonding character of MOs in the parent molecule. These approaches show an essentially non‐bonding character of the σ central C−C interaction in propellane. Within the MO theory, this bonding is thus only due to π‐type MOs (also called “banana” MOs or “bridge” MOs) and its total energy is evaluated to approximately 50 kcal mol−1. In bicyclopentane, despite a strong σ‐type repulsion, a weak bonding (15–20 kcal mol−1) exists between both central C−C bonds, also due to π‐type interactions, though no bond is present in the Lewis structure. Overall, the so‐called “inverted” bond, as resulting from a σ overlap of the two sp hybrids by their smaller lobes, appears highly questionable.

It is shown how use of valence bond structures of the type for a 5-electron 3-centre bonding unit can help to provide compact valence bond representations of the electronic structures for NO2, N2O4, excited O3 and SO2, SO3, SO2Y2 (with Y = OH−, O− or F), ClO2, SO2−, S2O42−, [Cu(H2O)6]2+ and an N H N linkage. Without atomic valence shell expansions, many of these structures provide electronic interpretations for hypervalent 19th century type bond diagrams. As an example, it is shown that the 5-electron 3-centre valence bond structure for N2O4 (which is equivalent to resonance between 144 canonical Lewis structures) provides an electronic interpretation for Frankland’s 1866 bond diagram with pentavalent nitrogen atoms.

An explanation is presented for the unusual conformations of some molecules that contain the C═C-C═C core, namely, butadienes, biphenyls, and styrenes. Small substituents often induce a synclinal conformation, which brings the substituents into close proximity, and sometimes, there is no anticlinal minimum at all. This would not be predicted from steric repulsion arguments nor would it be expected that atoms that are nonbonded in a Lewis structure would approach closer than the sum of their van der Waals radii. Atomic energies calculated according to the quantum theory of atoms in molecules (QTAIM) do not show a consistent pattern for these structurally similar molecules, nor are intersubstituent bond paths consistently found, nor favorable diatomic interaction energies calculated using the interacting quantum atoms (IQA) scheme. Instead, the synclinal conformations are found to be driven by the attraction energy of the electron distribution of the carbon atoms and the nuclei of the molecule.

There is a great variety of bond analysis tools that aim to extract information on the bonding situation from the molecular wavefunction. Because none of these can fully describe bonding in all of its complexity, it is necessary to regard a balanced selection of complementary analysis methods to obtain a reliable chemical conclusion. This is, however, not a feasible approach in most studies because it is a time-consuming procedure. Therefore, we provide the first comprehensive comparison of modern bonding analysis methods to reveal their informative value. The element-oxygen bond of neutral Hn XOH model compounds (X=Li, Be, B, C, N, O, F, Na, Mg, Al, Si, P, S, Cl) is investigated with a selection of different bond analysis tools, which may be assigned into three different categories: i) real space bonding indicators (quantum theory of atoms in molecules (QTAIM), the electron localizability indicator (ELI-D), and the Raub-Jansen index), ii) orbital-based descriptors (natural bond orbitals (NBO), natural resonance theory (NRT), and valence bond (VB) calculations), and iii) energy analysis methods (energy decomposition analysis (EDA) and the Q-analysis). Besides gaining a deep insight into the nature of the element-oxygen bond across the periodic table, this systematic investigation allows us to get an impression on how well these tools complement each other. Ionic, highly polarized, polarized covalent, and charge-shift bonds are discerned from each other.

A stand-alone polymer unit, centered on a tie-dye activity, was used to introduce the concept of polymers and the process of polymerization to three different groups of beginning-level chemistry students. This polymer unit consists of three parts. First, each student used a molecular model kit to construct the monomer units, β-d-glucose. Students combined these monomer units to learn the polymerization process and visualize the formation of the cellulose polymer in conjunction with writing the balanced equation for the polymerization reaction. Second, students tie-dyed cotton t-shirts made of cellulose fibers using dyes purchased from local craft stores. Finally, a primary literature exercise engaged students in understanding the chemical reactions that bind the cellulose fibers to the dye molecules permanently, producing a colorfast tie-dyed t-shirt. This stand-alone polymer unit can be used in any beginning college or high school chemistry course to introduce students to polymer chemistry and the proces...

This article dwells on the nature of “inverted bonds”, which make reference to the σ interaction between two s-p hybrids by their smaller lobes, and their presence in [1.1.1]propellane 1. Firstly we study H 3 C-C models of C-C bonds with frozen HCC angles reproducing the constraints of various degrees of “inversion”. Secondly, the molecular orbital (MO) properties of [1.1.1]propellane 1 and [1.1.1]bicyclopentane 2 are analyzed with the help of orbital forces as a criterion of bonding/antibonding character and as a basis to evaluate bond energies. Triplet and cationic state of 1 species are also considered to confirm the bonding/antibonding character of MOs in the parent molecule. These approaches show an essentially non-bonding character of the σ central CC interaction in propellane. Within MO theory, this bonding is thus only due to π-type MOs (also called ‘banana’ MOs or ‘bridge’ MOs) and its total energy is evaluated to ca. 50 kcal/mol. In bicyclopentane 2, despite a strong σ-type repulsion, a weak bonding (15-20 kcal/mol) exists between both central CC, also due to π-type interactions, though no bond is present in the Lewis structure. Overall, the so-called ‘inverted’ bond, as resulting from a σ overlap of the two s-p hybrids by their smaller lobes, appears highly questionable.

A simple manipulative resource, Atomic Tiles, is described for scaffolding the learning of Lewis structures without using algorithmic, rule-based methods of drawing. Students use Atomic Tiles to (1) create models of bonding that lead to drawing Lewis structures, (2) use the structures they create to infer patterns required for rational structures and common organic functional groups, (3) translate between Lewis structures and molecular models, and (4) use molecular models to identify isomers.

Consideration is given to (electronically) hypervalent increased-valence structures, which possess 2c–1e bonds, fractional 2c–2e bonds, and usually normal 2c–2e bonds., Consideration is given to (electronically) hypervalent increased-valence structures, which possess 2c–1e bonds, fractional 2c–2e bonds, and usually normal 2c–2e bonds. For singlet-spin electron-rich systems, increased-valence structures, with Heitler–London 2c–2e bond wavefunctions, are equivalent to resonance between non-hypervalent Kekulé and Dewar (or singlet diradical) type Lewis structures. Dewar structures are not considered in the Chem. Sci. 2015, 6, 6614 Edge article on hypervalency. Using one-electron delocalizations from lone-pair atomic orbitals into separate bonding molecular orbitals, increased-valence structures for PCl5, O3, SO42–, NO3–, N2O4 and SN2 reactions are derived from the Edge-article's Kekulé-type Lewis structures, and compared with the Edge article's hypervalent structures with 2c–2e bonds. It is also shown that Durrant's method to determine the γ parameter for XAY-type systems that possess a symmetrical 3c–4e bonding unit is related to the A-atom charge density.

A thermally robust triple-decker complex [bis(η5-pentamethylcyclopentadienyltitanium)-μ-(η4:η4-1,2,4,5-tetrakis(trimethylsilyl)cyclohexa-1-4-diene-3,6-diyl)] (3) was obtained in 4% yield by thermolysing Cp*TiMe3 in the presence bis(trimethylsilyl)acetylene (BTMSA). The solid-state structure of centrosymmetric 3 features rather long all C–C bonds in the nearly planar bridging ligand (1.4720(14)–1.4896(15) A) and a short distance of its least-square plane to the titanium atoms (1.7381(5) A). Computational results revealed the bonding of the central ligand to be accomplished through back-bonding of its two CC bonds and through the simultaneous generation of two σ-Ti–C(H) bonds. Based on CASSCF and CASPT2 results, the molecule acquires several electronic configurations simultaneously, which hinders its representation by one single Lewis structure. Apart from being coordinated to the central ligand, the metal atoms are involved in a direct Ti–Ti bonding by the formation of one σ- and two δ-bonds between them. The bond order of this Ti–Ti overlap shows only a slight decrease upon electronic excitation. The presence of ionic contribution to the bonding of the central ligand is manifested by the charge −1.4e summed on the carbon atoms of the bridging ring. Based on computational results, the spin multiplicity of the ground state is singlet, while the first low lying excitation state is triplet. This is in agreement with the absence of EPR signal in either toluene solution and glass, and with slight downfield shifts of broadened 1H NMR signals of SiMe3 and Cp* methyl groups observed with increasing temperature.

An embedding method for a one-electron reduced density matrix (1-RDM) is proposed. It is based on partitioning of 1-RDM into domains and describing each domain in the effective potential of the other ones. To assure N-representability of the total 1-RDM N-representability and strong-orthogonality conditions are imposed on the domains. The total energy is given as a sum of single-domain energies and domain-domain electron interaction contributions. Higher than two-body inter-domain interaction terms are neglected. The two-body correlation terms are approximated by deriving inter-domain correlation from couplings of density fluctuations of two domains at a time. Unlike in most density embedding methods kinetic energy is treated exactly and it is not required that densities pertaining to the domains are only weakly overlapping. We propose to treat each domain by a corrected perfect-pairing functional. On a few examples it is shown that the embedding reduced density matrix functional method (ERDMF) yields excellent results for molecules that are well described by a single Lewis structure even if strong static intra-domain or dynamic inter-domain correlation effects must be accounted for.

Chemical bonding is an important topic in chemistry. The chemistry students have to expert on the topic. Due to the importance of chemical bonding topic, qualitative descriptive research was conducted to analyze the students' errors in solving the chemical bond problems. The participants were 38 students of the third semester of Chemistry Education Department in one of the universities in Aceh, Indonesia. The result showed that the percentage of the students’ errors in solving the chemical bonds problems was high with an average of 69.08%, it means that the students’ ability in understanding chemical bond concepts is still low. This research concludes two things. First, the types of errors in solving chemical bonds problems are the students’ error in understanding the basic concepts of chemical bonds, lack of understanding of the concept how an element has a stable electron configuration, the mechanism of the difference in the forming of ion bonds and covalent bonds, the influence of geometric figures, the resultant molecular dipole moments, the number of electronegativity to molecular pollutants, the writing of a Lewis structure for molecules and polyatomic ions, the determining of the formal contents of each atom of a Lewis molecular structure, and the drawing of the Lewis structure. Second, the causes of students’ errors have forgotten the materials and also pre-requisite materials and inaccurate in reading and solving the problems. These results need to be considered in the learning process in order to improve students’ ability in solving chemical bond problems.

The activation of CO2 by chloride-tagged first-row transition metal anions [ClM]− (M = Sc–Zn), was examined by mass spectrometry, quantum chemical calculations, and statistical analysis. The direct formation of [ClM(CO2)]− complexes was demonstrated in the reaction between [ClM]− and neutral CO2. In addition, the reverse reaction was investigated by energy-variable collisionally induced dissociation (CID) of the corresponding [ClM(CO2)]− anions generated in-source. Five different mono- and bi-dentate binding motifs present in the ion/CO2 complexes were identified by quantum chemical calculations and the relative stability of each of these isomers was established and analyzed for all first-row transition metals based on the experimental and theoretical ion/molecule binding energies. It was found that the early first row transition metals form strong covalent bonds with the neutral CO2 molecule, while the late ones and in particular copper and zinc are weakly bonded. Using simple valence bond Lewis diagrams, the different binding motifs and their relative stabilities across the first row were described using multi-configurational self consistent field (MSCSCF) wavefunctions in a quantitative manner based on the electronic structure of the individual metals. This analysis provides an explanation for the change of the most favorite bonding motif of the transition metals with CO2 along the 1st transition metal row. The nature of the activated CO2 complex and the relationship between its stability and other structural and spectral properties was also analyzed by Principal Component Analysis (PCA) and artificial neural networks.

Author(s): Halbach, Robert L; Nocton, Gregory; Booth, Corwin H; Maron, Laurent; Andersen, Richard A | Abstract: Page 7295. In the Summary section of the article, the text should read "The energy of the closed-shell singlet that corresponds to the traditional Lewis structures used to represent these molecules is 400-700 cm-1 above the openshell triplet and singlet." instead of "The energy of the openshell singlet that corresponds to the traditional Lewis structures used to represent these molecules is 400-700 cm-1 above the open-shell triplet."

Lewisove strukture ili elektronske strukturne formule prikazuju na koji su način atomi u molekuli međusobno povezani te raspored valentnih elektrona svih atoma u molekuli. Ako se pored pravila za njihovo sastavljanje primijene još i VSEPR-teorija, te pojmovi kao što su formalni naboj, rezonancija i hibridizacija, dobije se prikaz koji vrlo dobro opisuje i građu i oblik molekule. Ovo djelo je dano na korištenje pod licencom Creative Commons Imenovanje 4.0 međunarodna., Lewis structures or Lewis dot structures show the way in which the atoms in the molecule are linked to each other, as well as the arrangement of valent electrons of all atoms in the molecule. Applying simple rules, which are more in accordance with the molecular orbital theory, the dot structures even of complicated molecules can be written with ease. If beside those rules we apply the VSEPR theory, we can deduce the structure and the shape of the molecule, and from the structure of the molecule, we can deduce the hybridization of the atoms. When we calculate the formal charges, from their values we can conclude if the structure is stable or more stable resonance structures exist. If the molecule in consideration contains atoms from the third period and on, in that case, more stable structures can be obtained by the formation of multiple bonds and expanded valence shell. In the present paper, the application of the above mentioned is shown step by step on several examples. This work is licensed under a Creative Commons Attribution 4.0 International License.

Coordination of a B2 fragment by two σ-donor ligands could lead to complexes with a formal triple B≡B bond L→B≡B←L. Formation of L-B σ-bond leads to excess electrons around the B2 central fragment. A subsequent direct π-charge transfer from B≡B moiety to the ligands L is a necessary condition for incorporation of BB fragment to the conjugated LBBL system. Quantum-chemical calculations (DFT, CCD, and CAS) show that the stabilization of a linear L-BB-L structure is possible but is accompanied by lowering of the B-B bond order. The ammonia-boryne structure H3N→BB←NH3, which is studied as a model system, shows a perfect triple BB fragment relative to other LBBL adducts. The comparison of the conjugation between the B2 fragment and two types of σ-donor ligands, with or without π-back-donation, provides an indication of the character of BB multiple bond. Three studied boryne molecules are calculated to have a high barrier for dissociation XXBBXX→XX + BBXX, yet a high reactivity of these compounds is indicated by the low-lying excited states of borynes (ΔES0-S1S0-S1∼2.6 eV is calculated for the room temperature stable bis-NHC-boryne.

Theoretical studies on Lewis acid-base behavior of hypervalent halogen fluorides, F3X and F5X (X = Cl, Br, I) have been instrumental in guiding this work. We have also examined whether the hole-lump concept explains the formation of the F5XCO complexes. Calculations of proton affinities (PA) and gas-phase basicity (GB) on hypervalent halogen fluorides show that F3X and F5X molecules can act as Lewis bases in gas phase. Moreover, theoretical calculations indicate that F3X and F5X molecules can act as Lewis acids forming stable complexes with a Lewis base as CO. The quantum theory of atoms in molecules (QTAIM) shows that the electrostatic interaction between the lone pair of the Lewis base (CO) and nucleus of the hypervalent halogen atom (X) plays a key role in stabilizing and determining the optimal geometry of the F5XCO complexes, as in conventional XBs. The localized molecular orbital energy decomposition analysis (LMOEDA) reveals that electrostatic component plays an important role in the stability of the FnXCO complexes. Fil: Buralli, Gabriel Jesús. Consejo Nacional de Investigaciones Científicas y Técnicas. Centro Científico Tecnológico Conicet - Nordeste. Instituto de Química Basica y Aplicada del Nordeste Argentino. Universidad Nacional del Nordeste. Facultad de Cs.exactas Naturales y Agrimensura. Instituto de Química Basica y Aplicada del Nordeste Argentino; Argentina. Universidad Nacional del Nordeste. Facultad de Ciencias Exactas y Naturales y Agrimensura. Departamento de Química. Laboratorio de Estructura Molecular y Propiedades; Argentina Fil: Duarte, Darío Jorge Roberto. Universidad Nacional del Nordeste. Facultad de Ciencias Exactas y Naturales y Agrimensura. Departamento de Química. Laboratorio de Estructura Molecular y Propiedades; Argentina. Consejo Nacional de Investigaciones Científicas y Técnicas. Centro Científico Tecnológico Conicet - Nordeste. Instituto de Química Basica y Aplicada del Nordeste Argentino. Universidad Nacional del Nordeste. Facultad de Cs.exactas Naturales y Agrimensura. Instituto de Química Basica y Aplicada del Nordeste Argentino; Argentina Fil: Sosa, Gladis Laura. Consejo Nacional de Investigaciones Científicas y Técnicas. Centro Científico Tecnológico Conicet - Nordeste. Instituto de Química Basica y Aplicada del Nordeste Argentino. Universidad Nacional del Nordeste. Facultad de Cs.exactas Naturales y Agrimensura. Instituto de Química Basica y Aplicada del Nordeste Argentino; Argentina. Universidad Tecnológica Nacional. Facultad Regional Resistencia. Departamento de Ingeniería Química. Laboratorio de Química Teórica y Experimental; Argentina Fil: Peruchena, Nelida Maria. Universidad Nacional del Nordeste. Facultad de Ciencias Exactas y Naturales y Agrimensura. Departamento de Química. Laboratorio de Estructura Molecular y Propiedades; Argentina

Topological analysis of the Electron Localisation Function (ELF) within the framework of Quantum Chemical Topology (QCT) has been applied to study the nature of the boron–oxygen bonds. A series of 16 compounds has been chosen, with the experimental B O bond length in the range of 1.481 A (B O)–1.179 A (B O). Topological results obtained for the DFT(M062X), DFT(B3LYP), MP2 and CCSD(T) optimised geometrical structures show that all the boron–oxygen bonds in the investigated compounds are described by the disynaptic bonding basin, V(B,O). All these bonds have a covalent-polarised character. The mean electron population of V(B,O) varies from 1.6e (B(OH) 4 − ) to about 3.5e (HN CH CH CH NH B O). The polarity index values, p BO, lie between 0.77 (ClBO) and 0.89 (H 2 BOCH 3 ), thus all boron–oxygen bonds are essentially polarised by the oxygen atom. According to the Lewis formula, four types of the bonds have been recognised. These are: a single bond with a mixture of the ionic hybrid (B O, B + O − ), a single bond (B O), a single bond with a small contribution of the dative O→B bond (B O) and a single bond with a large contribution of the dative O→B bond (depleted B O bonds). There is a clear distinction between a group of 11 molecules chosen for this study, with the basin population value of the boron–oxygen bond between 1.6e and 2.4e, and the HB O, FB O, ClB O, HN CH CH CH NH B O and trans-[(Me3P)2BrPt(B O)] molecules that exhibit the basin population in the range: 3.3e–3.5e. The second group was postulated to have a triple bond, B O, but this statement has not been confirmed by our research.

We call iodabenzene a cyclic (CH)5I molecule. A planar iodabenzene would have 8 π electrons, a situation best avoided by an out-of-plane distortion to a bird-like geometry. The electronic structure and charge distribution of this molecule resemble those of Meisenheimer complexes, derivatives of (CH)5CH2–. A similar substitution strategy, of π-acceptors in ortho and para positions, works in both cases to planarize and stabilize such derivatives. Some 40 kcal/mol (73 kcal/mol for the unsubstituted case) below the bird, a classical 5-iodocyclopentadiene structure awaits, reached through a bicyclic transition state. The calculated activation barrier for the highly exothermic reaction to a classical Lewis structure nevertheless make us optimistic about the chances of detecting and even isolating the bird isomer.

The borane chemistry has continued to expand since the beginning of the twentieth century. It has deeply penetrated into the other fields such as organic chemistry, inorganic chemistry, organometallic chemistry, catalysis, and medicinal fields. Yet some of the average chemists or chemists at an undergraduate or even postgraduate levels find it a challenge to explain the structures of simple boranes. In this article, we present an empirical formula that can be utilized to deduce geometrical shapes of simplesboranes. The formula is so versatile as to be applied to deduce the Lewis structures of simple molecules and ions, inorganic clusters of main group elements, the carbonyl clusters of transition metal complexes as well as carboranes symmetries and hydrocarbons. In the case of hydrocarbons, it is useful in generating isomers of a given hydrocarbon. It is hoped that the formula will be extended to explain the structures or shapes and symmetries of large clusters and to systems which may not obey the octet or eighteen electron rule.

We design theoretically small molecule candidates for singlet fission chromophores, aiming to achieve a balance between sufficient diradical character and kinetic persistence. We develop a perturbation strategy based on the captodative effect to introduce diradical character into small π-systems. Specifically, this can be accomplished by replacing pairs of not necessarily adjacent C atoms with isoelectronic and isosteric pairs of B and N atoms. Three rules of thumb emerge from our studies to aid further design: (i) Lewis structures provide insight into likely diradical character; (ii) formal radical centers of the diradical must be well-separated; (iii) stabilization of radical centers by a donor (N) and an acceptor (B) is essential. Following the rules, we propose candidate molecules. Employing reliable multireference calculations for excited states, we identify three likely candidate molecules for SF chromophores. These include a benzene, a napthalene, and an azulene, where four C atoms are replaced by a pair of B and a pair of N atoms.

Covalent bond dissociation is examined by three geminal-based theories. One approach (antisymmetrised product of strongly orthogonal geminals) assumes purely singlet geminals, while two others operate with mixtures of singlets and triplets (cf. restricted–unrestricted strongly orthogonal singlet-type geminals, hereafter RUSSG). The RUSSG is the wavefunction in the second method. It is spin contaminated. The spin contamination of RUSSG is projected out in a spin-purification step in the third method. Description of local (i.e. atomic) spin by geminal-based theories is examined. Prototype systems showing the deficiency of singlet coupling are taken as test cases. We find that the local spin of equilibrium structures is correctly described by purely singlet geminals. Triplet geminals are shown to be essential for the description of local spin when dissociating multiple bonds, or switching between two Lewis structures of the same molecule.

Quantum calculations at the DFT-D3/def2-TZVPD level of theory have been used to examine complexes between O2YBr (Y═N, P, and As) molecules and several Lewis bases, that is, NH3, H2O, and HF. The interactions of the lone pair of the ammonia, water, and hydrogen fluoride with the σ-hole and π-hole of O2YBr molecules have been considered. In general, the complexes where the Lewis base lone pair interacts with the π-hole are more favorable than those with σ-hole. The nature of the interactions has been characterized with the Bader theory of atoms in molecules (AIM). We have also studied the ability of trifluoronitromethane and nitromethane to interact with anions using their π-hole along with an analysis the Cambridge Structural Database. We have found a large number of hits that provide strong experimental support for ability of the nitryl (-NO2) group to interact with anions and Lewis bases. In some X-ray structures, the π-hole interaction is crucial in the crystal packing and has a strong influence in the solid state architecture of the complexes. Finally, due to the relevance in atmospheric chemistry, we have studied noncovalent σ/π-hole complexes of nitryl bromide with ozone.

This work aims to contribute to the discussion about the suitability of bond paths and bond-critical points as indicators of chemical bonding defined within the theoretical framework of the quantum theory of atoms in molecules. For this purpose, we consider the temporal evolution of the molecular structure of [Fe{C(CH2 )3 }(CO)3 ] throughout Born-Oppenheimer molecular dynamics (BOMD), which illustrates the changing behaviour of the molecular graph (MG) of an electronic system. Several MGs with significant lifespans are observed across the BOMD simulations. The bond paths between the trimethylenemethane and the metallic core are uninterruptedly formed and broken. This situation is reminiscent of a "hopping" ligand over the iron atom. The molecular graph wherein the bonding between trimethylenemethane and the iron atom takes place only by means of the tertiary carbon atom has the longest lifespan of all the considered structures, which is consistent with the MG found by X-ray diffraction experiments and quantum chemical calculations. In contrast, the η(4) complex predicted by molecular-orbital theory has an extremely brief lifetime. The lifespan of different molecular structures is related to bond descriptors on the basis of the topology of the electron density such as the ellipticities at the FeCH2 bond-critical points and electron delocalisation indices. This work also proposes the concept of a dynamic molecular graph composed of the different structures found throughout the BOMD trajectories in analogy to a resonance hybrid of Lewis structures. It is our hope that the notion of dynamic molecular graphs will prove useful in the discussion of electronic systems, in particular for those in which analysis on the basis of static structures leads to controversial conclusions.

In organic chemistry, as defined by Abegg, Kossel, Lewis and Langmuir, compounds are normally represented using structural formulas called Lewis structures. In these structures, the octet rule is used to define the number of covalent bonds that each atom forms with its neighbors and multiple bonds are frequent. Lewis’ octet rule has unfortunately shown limitations very early when applied to non-organic compounds: most of them remain incompatible with the “rule of eight” and location of charges is uncertain. In an attempt to unify structural formulas of octet and non-octet molecules or single-charge ions, an even-odd rule was recently proposed, together with a procedure to locate charge precisely. This even-odd rule has introduced a charge-dependent effective-valence number calculated for each atom. With this number and the number of covalent bonds of each element, two even numbers are calculated. These numbers are both used to understand and draw structuralformulas of single-covalent-bonded compounds. In the present paper, a procedure is proposed to adjust structural formulas of compounds that are commonly represented with multiple bonds. In order to keep them compatible with the even-odd rule, they will be represented using only single covalent bonds. The procedure will then describe the consequences of bond simplification on charges locations. The newly obtained representations are compared to their conventional structural formulas, i.e. single-bond representation vs. multiple-bond structures. Throughout the comparison process, charges are precisely located and assigned to specific atoms. After discussion of particular cases of compounds, the paper finally concludes that a rule limiting representations of multiplecovalent bonds to single covalent bonds, seems to be suitable for numerous known compounds.

The concept of covalent bonding is characterized by an interconnected knowledge framework based on Lewis and quantum models of atoms and molecules. Several research studies have shown that students at all levels of chemistry learning find the quantum model to be one of the most difficult subjects to understand. We have tried in this paper to analyze the extent to which Algerian students, at the end of their training, have integrated the covalent bonding theories based on the quantum model of atom theory and are able to interpret Lewis structures using the quantum model. The analysis of the responses to a written questionnaire showed that this integration was not achieved by our students and that they are not able to correctly describe covalent bonds in a Lewis structure using the concepts of the quantum model. They have a “quantum box” conception of atomic or hybrid orbitals. This conception acts as a “pedagogical learning impediment” to the integration of the geometrical representations of atomic and hybrid orbitals, the conditions of their overlapping to give bonds and consequently the description of covalent bonds using the quantum model. So, the students use an alternative conceptual framework based on the use of Lewis model paired valence electrons to form covalent bonds that we have named the: “electrons pair framework”. Furthermore, the students restricted the denomination of a covalent bond to the sharing of one electron (either s or p but not spn) from each atom to give one “electron pair σ”, and thought that σ bonds are only formed in single bonds.

Lewis developed a 2D-representation of molecules, charged or uncharged, known as structural formula, and stated the criteria to draw it. At the time, the vast majority of known molecules followed the octet-rule, one of Lewis’s criteria. The same method was however rapidly applied to represent compounds that do not follow the octet-rule, i.e. compounds for which some of the composing atoms have greater or less than eight electrons in their valence shell. In a previous paper, an even-odd rule was proposed and shown to apply to both types of uncharged molecules. In the present paper, the even-odd rule is extended with the objective to encompass all single-bonded ions in one group: Lewis’s ions, hypo- and hypervalent ions. The base of the even-odd representation is compatible with Lewis’s diagram. Additionally, each atom is subscripted with an even number calculated by adding the valence number, the number of covalent bonds of the element, and its electrical charge. This paper describes how to calculate the latter number and in doing so, how charge and electron-pairs can actually be precisely localized. Using ions known to be compatible with Lewis’s rule of eight, the even-odd rule is compared with the former. The even-odd rule is then applied to ions known as hypo- or hypervalent. An interesting side effect of the presented rule is that charge and electron-pairs are unambiguously assigned to one of the atoms composing the single-charged ion. Ions that follow the octet rule and ions that do not, are thus reconciled in one group called “electron-paired ions” due to the absence of unpaired electrons. A future paper will focus on the connection between the even-odd rule and molecules or ions having multiple bonds.

Viewpoint pubs.acs.org/JPCA 4th International Conference on Chemical Bonding T that could be made during the ICCB series is the synergy between the very diverse subdisciplines in the field of chemical bonding. One example, of the many fascinating contributions, is the bonding between nanoparticles that behave as very large atoms that are held together by DNA linkers playing the role of chemical bonds (Park, S. Y.; Lytton-Jean, A. K.; Lee, B.; Weigand, S.; Schatz, G. C.; Mirkin, C. A. DNA-programmable nanoparticle crystallization. Nature 2008, 451, 553−556). This topic is relevant for designing functional molecular materials but it also provides opportunities to probe fundamental bonding concepts from the realm of small molecules composed of regular atoms, such as hypervalence. Although the fundamental core of the meeting stays the same, the focus on different kinds of exciting systems shifts from year to year. In 2016, the prominent themes included chemical bonding at interfaces, in clusters, nanoparticles, 2D materials, and materials with outstanding mechanical proper- ties. Here, we will highlight some of these themes. he fourth International Conference on Chemical Bonding (ICCB) was held in Kauai Island, Hawaii, from July 14th to 26th, 2016. Sixty speakers from 15 different countries around the world gathered on this beautiful island for a multi- disciplinary conference organized by Profs. Alexander Boldyrev and Anastassia Alexandrova, and sponsored by SCM - Software for Chemistry & Materials, The Journal of Chemical Theory and Computation, and the Chemistry & Biochemistry Departments of Utah State University and the University of California, Los Angeles. This brief viewpoint article summarizes the themes and objective of the ICCB Conference series. A full list of the presentations with their corresponding abstracts can be downloaded from the Conference web page: http://www. chem.ucla.edu/~ana/ICCB2016.html The link between electronic structure and properties is, of course, central to chemistry. Nowadays it is possible to perform first principles calculations or electronic spectroscopy experi- ments on chemical systems, to probe their electronic structure and derive sound reasons for their observed properties. However, instead of doing this for each and every molecule we encounter, we instead generalize our observations into the language of chemical bonding, a set of qualitative concepts that allows chemists to converse efficiently with one another, teach students, and create and pass ideas quickly. It would take a good chemist only seconds to go from drawing a Lewis structure of a molecule to stating its properties with fair accuracy. The theory of chemical bonding ultimately gives chemists a huge advantage in understanding and designing matter. As an undisclosed UCLA fellow once said: “Without chemical bonding, we all still would be solving the Schro dinger equation for H 2 on cave walls; we would have no polymers, no modern drugs, and no detergents, to name a few”. Despite the apparent utility of the language of chemical bonding, it is incomplete: for many interesting systems, such as alloys, catalytic surfaces, nanoparticles, and clusters (including such long-known clusters as Fe−S and FeMoCo in biology) the theory is just not yet developed. In other words, we are not ready to draw an appropriate analogue of a Lewis structure, and from there, start speaking about properties. It is therefore of paramount interest to build this theory and start using it in the design of the cutting-edge materials for the future. This is why scientists gather on the Pacific island of Kauai, in addition to enjoying its legendary beauty. The ICCB series is devoted to bringing together theorists and experimentalists, and to putting the fuzzy concept of the chemical bond onto firmer grounds. From the theoretical perspective, developing computational tools to make chemical bonds quantitatively accurate is the fundamental key to reach the ultimate objective of allowing the discovery of new molecules, functional materials, and homogeneous and heterogeneous catalysts and understanding and predicting their properties in an intuitive and quick manner. Experimen- talists willing to take advantage of these tools come to unveil novel chemical bonding patterns and observable/measurable properties in all kinds of materials. An important observation © 2016 American Chemical Society FUNDAMENTALS Chemical bonding lies at the heart of virtually every model used in chemistry to classify and even predict new compounds. Much of the classification relies on the analysis of the nature of the chemical bonds in these new compounds. However, given that the chemical bond is not a uniquely defined object from the perspective of quantum mechanics, its nature remains to some extent elusive especially in novel classes of compounds where the H 2 paradigm does not apply so easily anymore. The key importance of a model for the nature of the chemical bond is clear from the attention devoted to such a model in several lectures. Three groups of methods that go back to the earliest days of quantum mechanical manifestations of the chemical bond, again played a major role in the 2016 conference. Miroslav Kohout (Max-Planck-Institut fu r Chemi- sche Physik fester Stoffe, Dresden, Germany) focused on the direct manifestation of chemical bonding in position space through sets of descriptors that can be used for the classification of bonding. Such an approach goes beyond the up-to-now more often used orbital approach and also beyond the simple depiction of electron densities. Cle mence Corminboeuf (EPFL Lausanne, Switzerland) showcased the Density Overlap Region Indicator (DORI) as a position space based descriptor that reveals different types of bonding and noncovalent interactions. In a second line of approaches, the focus lies on the energetic aspect of chemical bonding. Why does the energy lower when atoms link up on the basis of the variation principle? A series of presentations at ICCB2016 focused on energy decomposition analyses (EDA). A feature of these methods is that the difference between the energy of a structure and that of its composing fragments is decomposed in chemically relevant parts by stepwise releasing constraints going from the fragments to the final structure. Klaus Ruedenberg and Mark S. Gordon (Iowa State University, Ames, IA, USA) presented Published: November 18, 2016 DOI: 10.1021/acs.jpca.6b11179 J. Phys. Chem. A 2016, 120, 9353−9356

Probing the electron density transfers during a chemical reaction can provide important insights, making possible to understand and control chemical reactions. This aim has required extensions of the relationships between the traditional chemical concepts and the quantum mechanical ones. The present work examines the detailed chemical insights that have been generated through 100 years of work worldwide on G. N. Lewis's ground breaking paper on The Atom and the Molecule (Lewis, G. N. The Atom and the Molecule, J. Am. Chem. Soc. 1916, 38, 762–785), with a focus on how the determination of reaction mechanisms can be reached applying the bonding evolution theory (BET), emphasizing how curly arrows meet electron density transfers in chemical reaction mechanisms and how the Lewis structure can be recovered. BET that combines the topological analysis of the electron localization function (ELF) and Thom's catastrophe theory (CT) provides a powerful tool providing insight into molecular mechanisms of chemical rearrangements. In agreement with physical laws and quantum theoretical insights, BET can be considered as an appropriate tool to tackle chemical reactivity with a wide range of possible applications. Likewise, the present approach retrieves the classical curly arrows used to describe the rearrangements of chemical bonds for a given reaction mechanism, providing detailed physical grounds for this type of representation. The ideas underlying the valence-shell-electron pair-repulsion (VSEPR) model applied to non-equilibrium geometries provide simple chemical explanations of density transfers. For a given geometry around a central atom, the arrangement of the electronic domain may comply or not with the VSEPR rules according with the valence shell population of the considered atom. A deformation yields arrangements which are either VSEPR defective (at least a domain is missing to match the VSEPR arrangement corresponding to the geometry of the ligands), VSEPR compliant or pseudo VSEPR when the position of bonding and non-bonding domains are interchanged. VSEPR defective arrangements increase the electrophilic character of the site whereas the VSEPR compliant arrangements anticipate the formation of a new covalent bond. The frequencies of the normal modes which account for the reaction coordinate provide additional information on the succession of the density transfers. This simple model is shown to yield results in very good agreement with those obtained by BET.

The electron-pair bonding model that was introduced by Gilbert Lewis 100 years ago is discussed in the light of modern quantum chemical methods for analysing the electronic structures of some simple molecules. It is argued that Lewis structures in conjunction with accurate quantum chemical calculations are still very useful for the description of chemical bonding. The emphasis lies on the difference between electron-sharing bonds A–B and dative bonds A → B which were suggested by Lewis as a general definition for acids and bases. The electron-pair model, if combined with quantum chemical calculations, remains a powerful guide for the search of new molecules and for understanding molecular structures.

A discovery-based method is described for incorporating the concepts of IR and Raman spectroscopy into the general chemistry curriculum. Students use three sets of springs to model the properties of single, double, and triple covalent bonds. Then, Gaussian 03W molecular modeling software is used to illustrate the relationship between bond vibrations and both IR and Raman spectroscopy. Students view the characteristic vibrations of C–C, C═C, and C≡C bonds in propane, propene, and propyne molecules. Finally, students collect IR and Raman spectra of hexane, 1-hexene, and 1-heptyne. This series of activities is completed in the first semester of the general chemistry sequence following the development of Lewis structures and before the discussion of electronegativity and bond polarity.

Natural bond orbital (NBO) methods encompass a suite of algorithms that enable fundamental bonding concepts to be extracted from Hartree-Fock (HF), Density Functional Theory (DFT), and post-HF computations. NBO terminology and general mathematical formulations for atoms and polyatomic species are presented. NBO analyses of selected molecules that span the periodic table illustrate the deciphering of the molecular wavefunction in terms commonly understood by chemists: Lewis structures, charge, bond order, bond type, hybridization, resonance, donor–acceptor interactions, etc. Upcoming features in the NBO program address ongoing advances in ab initio computing technology and burgeoning demands of its user community by introducing major new methods, keywords, and electronic structure system/NBO communication enhancements. © 2011 John Wiley & Sons, Ltd.

Main factors responsible for the red to near infrared (NIR) absorption of the benzobis(1,2,5-thiadiazole) (BBT) derivatives have been investigated using high level computational studies. These molecules with NIR emission are of importance due to the recent reports of possible role in NIR organic light emitting diodes (OLED) development. The electronic structure has been determined using closed-shell/open-shell DFT methods (B3LYP and BHandHLYP). The wavefunction stabilities of these BBT derivatives have been tested. We notice that using the BHandHLYP functional, the wave function becomes instable though large HOMO–LUMO gaps (HLG) are obtained. On the other hand a stable wavefunction is obtained using the B3LYP functional but the HLG is small. The B3LYP HLG is in good agreement with the available experimental data. Nevertheless results from both functionals indicate a possible LUMO occupation/biradicaloid character (BRC). We calculated the BRC for all the molecules using different methods and observed that these molecules have a large BRC which correlates well with excitation energies. Larger the BRC smaller the excitation energy. Charge transfer based on Mulliken charges of both ground and excited state are obtained from high level SAC/SAC-CI studies. We carried out the VB studies of the unsubstituted BBT to predict the relative weights of the individual Lewis structures to the resonance picture. It is concluded that a more general description which include the zwitterionic and biradical structures are necessary for these molecules and not just simple donor–acceptor–donor (D–A–D) structures as suggested in the literature. Copyright © 2011 John Wiley & Sons, Ltd.

Epik is a computer program for predicting pK(a) values for drug-like molecules. Epik can use this capability in combination with technology for tautomerization to adjust the protonation state of small drug-like molecules to automatically generate one or more of the most probable forms for use in further molecular modeling studies. Many medicinal chemicals can exchange protons with their environment, resulting in various ionization and tautomeric states, collectively known as protonation states. The protonation state of a drug can affect its solubility and membrane permeability. In modeling, the protonation state of a ligand will also affect which conformations are predicted for the molecule, as well as predictions for binding modes and ligand affinities based upon protein-ligand interactions. Despite the importance of the protonation state, many databases of candidate molecules used in drug development do not store reliable information on the most probable protonation states. Epik is sufficiently rapid and accurate to process large databases of drug-like molecules to provide this information. Several new technologies are employed. Extensions to the well-established Hammett and Taft approaches are used for pK(a) prediction, namely, mesomer standardization, charge cancellation, and charge spreading to make the predicted results reflect the nature of the molecule itself rather just for the particular Lewis structure used on input. In addition, a new iterative technology for generating, ranking and culling the generated protonation states is employed.

The technique described in Chapter 12 for constructing “increased-valence” structures, namely that of delocalizing one or more non-bonding electrons of a standard Lewis structure into adjacent bonding orbitals, is quite general and easily applied. We shall now use this method to construct “increased-valence” structures for numerous molecular systems that involve 6-electron 5-centre, 8-electron 6-centre and longer N-centre bonding units, as well as for some molecules with 4-electron 3- centre and 6-electron 4-centre bonding units. In general, we shall find that only one or two “increased-valence” structures are required in order to make deductions concerning bond lengths that are in qualitative accord with the measured lengths, i.e. for the systems considered, “increased-valence” structures provide easily derived and economical representations of their electronic structures.

We are now ready to examine in detail the incorporation of the Pauling “3-electron bond” structure Ȧ · Ḃ into the valence-bond structures for electron-rich molecules that involve 4-electron 3-centre and 6-electron 4-centre bonding units. To do this, we may use any of three alternative methods. In this Chapter, we shall discuss one of them. It involves the spin-pairing of the unpaired-electron of Ȧ · Ḃ with the unpaired electron of either an atom Ẏ or a second Pauling “3-electron bond” structure Ċ · Ḋ.

Valence-bond structures with Pauling “3-electron bonds” between pairs of atoms may be written down for a number of triatomic radicals. Here we shall examine these types of structures for some radicals with either 17 or 19 valence-shell electrons. For these systems, it is necessary to write down two Pauling “3-electron bond” structures, which participate in resonance. The delocalized molecular orbital equivalent of this resonance involves the construction of three 3-centre molecular orbitals to accommodate five electrons; this is described in Section 6-4.

The B3LYP/cc-pvtz approach is used to obtain Kohn–Sham wavefunctions for NO2 and N2O4 and several related compounds. ELF and AIM analyses on these systems lead to a better understanding of the unusual character of the NN bond in N2O4 in which competing Lewis structures with one or two electrons on the nitrogen atom in NO2 lead to a very weak sigma covalent bond that receives essentially no bonding contribution from the pπ part of the molecule.

The linear TiCH molecule is studied by ab initio quantum mechanical calculations using quantitative basis sets and highly correlated computational methods. Potential energy curves along the Ti−CH coordinate have been computed to obtain a better understanding of molecular formation in eight low-lying states of the molecule. Total energies, dissociation energies (with respect to Ti+CH), equilibrium distances, and dipole moments are reported. Simple valence bond Lewis diagrams are used to interpret the nature of the bonding in all of the states studied.

The Electron Localisation Function, ELF pictorially visualises chemists' intuitive ideas of single and multiple bonds as well as non-bonding electron pairs in molecules. The power of the representation of chemical bonds via ELF is that on the one hand covalent, polar, and ionic bonds are distinguishable, and that on the other hand ELF can be calculated for molecules and solids. This enables us to transfer the ideas of chemical bonding from molecular to intermetallic compounds. Localised two-electron-two-centre bonds and lone pairs are present in solid-state valence compounds (Zintl phases) as expected by the 8-N rule. In solids, lone pairs are generally more contracted than in molecules due to 'lone-pair repulsion'. In intermetallic compounds localised electrons predominantly occur in the form of lone pairs. Lattice vibrations influence the strength of lone pair interactions and non-bonded interactions lead to an exchange of delocalised and localised electrons. Such a mechanism of local electron pair formation gives rise to ideas of a chemical view of the phenomenon of superconductivity in intermetallic compounds.

Because the presently used Lewis diagrams have their roots in static electron theories of the early twentieth century, this paper proposes substituting a double-arrow symbolism for the customary double dots. This means that students should be taught that a bond line (or “stick” as some call it) means a double arrow, not a double dot. Lone pairs are to be represented by double-curved arrows and later abbreviated to the old familiar double dot. There are obvious advantages to this new notation when atoms have different electronegativities. The double arrows can suggest the shifting of the bond pair equilibrium toward the most electronegative atom. The new notation is more compatible with VSEPR theory, and it seems to suggest the loci of largely imaginary MOs and AOs by the configuration of the arrows. With the help of some imagination, the new symbolism can clarify resonance in aromatic molecules and conjugated chains. The most radical suggestion is that resonance forms be abandoned in favor of conjugated blur bonds. The concept of blur bonds has the merit of being easily extendable to the explanation of metal bonding and electron conduction.