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10.5:

Valentiebindingstheorie

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Chemistry
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JoVE Core Chemistry
Valence Bond Theory

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Valentiebindingstheorie is een methode die wordt gebruikt om te begrijpen hoe de atomaire orbitalen van het kwantummechanische model kunnen overlappen om een covalente binding te produceren. Het veronderstelt dat bindingen worden gevormd wanneer atomaire interacties de algehele energie van het systeem verlagen. Denk aan de vorming van een waterstofmolecuul.Elk atoom heeft een enkel elektron in zijn 1s-orbitaal. Wanneer ze ver uit elkaar staan, ondervinden de waterstofatomen geen aantrekking of afstoting van elkaar, en wordt de energie van het systeem als nul beschouwd. Naarmate de atomen dichter bij elkaar komen, voelt elk elektron de aantrekkingskracht van de kern in het andere atoom.Tegelijkertijd stoten de elektronen elkaar af, net als de kernen. Als de attracties sterker zijn dan de afstotingen, neemt de energie van het systeem af naarmate de atomen elkaar naderen. De minimale potentiële energie wordt bereikt wanneer de elektron-elektron-en kern-kernafstotingen de aantrekkingskrachten tussen elektronen en kernen in evenwicht brengen.Voor waterstofmoleculen gebeurt dit wanneer de bindingslengte 74 picometer is. Op dit punt treedt een aanzienlijke overlap van de twee waterstof-1s-orbitalen op en vormt een covalente binding. De twee elektronen met tegengestelde spins worden aangetrokken door beide kernen en liggen in de ruimte die wordt gedeeld door beide atomaire orbitalen.Als de internucleaire afstand verder wordt verkleind, begint de energie te stijgen, voornamelijk als gevolg van de elektrostatische afstoting tussen de kernen. De valentiebindingstheorie stelt voor dat een chemische binding ontstaat door een overlap van gedeeltelijk gevulde atomaire orbitalen, inclusief die anders dan sferische s-orbitaal. In waterstoffluoride kunnen de halfgevulde 1s-orbitaal van waterstof en de halfgevulde 2p-orbitaal van fluor een wisselwerking hebben.De p-orbitaal ligt langs de internucleaire as en overlapt met de s-orbitaal van waterstof en vormt een binding. Wanneer een enkele binding wordt gevormd tussen twee niet-sferische orbitalen, zullen de twee orbitalen elkaar overlappen. De covalente binding in een fluormolecuul wordt gevormd door de overlap van twee halfgevulde p-orbitalen die zorgen voor maximale overlap en een sterkere binding.De covalente binding van het type gevormd door een head-to-head overlapping van atomaire orbitalen wordt een sigma-binding genoemd.

10.5:

Valentiebindingstheorie

Overview of Valence Bond Theory

Valence bond theory describes a covalent bond as the overlap of half-filled atomic orbitals (each containing a single electron) that yield a pair of electrons shared between the two bonded atoms. The orbitals on two different atoms overlap when a portion of one orbital and a portion of a second orbital occupy the same region of space. According to valence bond theory, a covalent bond results when two conditions are met: (1) an orbital on one atom overlaps an orbital on a second atom and (2) the single electrons in each orbital combine to form an electron pair. The mutual attraction between this negatively charged electron pair and the two atoms’ positively charged nuclei serves to physically link the two atoms through a force we define as a covalent bond. The strength of a covalent bond depends on the extent of overlap of the orbitals involved. Orbitals that overlap extensively form bonds that are stronger than those that have less overlap.

Effect of Orbital Overlap on the Energy of the System

The energy of the system depends on how much the orbitals overlap. In the case of hydrogen atoms, the sum of the energies of two hydrogen atoms changes as they approach each other. When the atoms are far apart, there is no overlap, and by convention, the sum of the energies is zero. As the atoms move together, their orbitals begin to overlap. Each electron begins to feel the attraction of the nucleus in the other atom. In addition, the electrons begin to repel each other, as do the nuclei. While the atoms are still widely separated, the attractions are slightly stronger than the repulsions, and the energy of the system decreases and a bond begins to form. As the atoms move closer together, the overlap increases, so the attraction of the nuclei for the electrons continues to increase as do the repulsions among electrons and between the nuclei. At some specific distance between the atoms, which varies depending on the atoms involved, the energy reaches its lowest (most stable) value. This optimum distance between the two bonded nuclei is the bond distance between the two atoms. The bond is stable because, at this point, the attractive and repulsive forces combine to create the lowest possible energy configuration. If the distance between the nuclei were to decrease further, the repulsions between nuclei and the repulsions as electrons are confined in closer proximity to each other would become stronger than the attractive forces. The energy of the system would then rise, resulting in the destabilization of the system. 

Bond Energy

The bond energy is the difference between the energy minimum, which occurs at the bond distance, and the energy of the two separated atoms. This is the quantity of energy released when the bond is formed. Conversely, the same amount of energy is required to break the bond. For a H2 molecule, at the bond distance of 74 pm, the system is 7.24 × 10−19 J lower in energy than the two separated hydrogen atoms. This may seem like a small number. However, we know from our earlier description of thermochemistry that bond energies are often discussed on a per-mole basis. For example, it requires 7.24 × 10−19 J to break one H–H bond, but it takes 4.36 × 105 J to break 1 mole of H–H bonds. 

Types of bonds

In addition to the distance between two orbitals, the orientation of orbitals also affects their overlap (other than for two s orbitals, which are spherically symmetric). Greater overlap is possible when orbitals are oriented such that they overlap on a direct line between the two nuclei. 

The overlap of two s orbitals (as in H2), the overlap of an s orbital and a p orbital (as in HCl), and the end-to-end overlap of two p orbitals (as in Cl2) all produce sigma bonds (σ bonds).

 A σ bond is a covalent bond in which the electron density is concentrated in the region along the internuclear axis; that is, a line between the nuclei would pass through the center of the overlap region. Single bonds in Lewis structures are described as σ bonds in valence bond theory.

A pi bond (π bond) is a type of covalent bond that results from the side-by-side overlap of two p orbitals. In a π bond, the regions of orbital overlap lie on opposite sides of the internuclear axis. Along the axis itself, there is a node, that is, a plane with no probability of finding an electron.

While all single bonds are σ bonds, multiple bonds consist of both σ and π bonds. As per the Lewis structures, O2 contains a double bond, and N2 contains a triple bond. The double bond consists of one σ bond and one π bond, and the triple bond consists of one σ bond and two π bonds. Between any two atoms, the first bond formed will always be a σ bond, but there can only be one σ bond in any one location. In any multiple bond, there will be one σ bond, and the remaining one or two bonds will be π bonds. With regards to bond energy, an average carbon-carbon single bond is 347 kJ/mol, while in a carbon-carbon double bond, the π bond increases the bond strength by 267 kJ/mol. Adding an additional π bond causes a further increase of 225 kJ/mol. We can see a similar pattern when we compare other σ and π bonds. Thus, each individual π bond is generally weaker than a corresponding σ bond between the same two atoms. In a σ bond, there is a greater degree of orbital overlap than in a π bond.

This text has been adapted from Openstax, Chemistry 2e, Section 8.1 Valence Bond Theory.