Valence Theory Shit Flashcards
Covalent bond forms when 2 conditions are met, what are they?
A covalent bond forms when two conditions are met: (1) an orbital on one atom overlaps with an orbital on a second atom, and (2) the single electrons in each orbital combine to create an electron pair. The mutual attraction between this negatively charged electron pair and the positively charged nuclei of the two atoms serves to physically link the atoms. This attraction creates a force that we define as a covalent bond. The strength of a covalent bond depends on the extent of overlap of the orbitals involved, with more extensive overlap resulting in stronger bonds.
Valence bond theory describes what?
Valence bond theory describes a covalent bond as the overlap of half-filled atomic orbitals, each containing a single electron. This overlap yields a pair of electrons shared between the two bonded atoms. We consider orbitals on two different atoms to overlap when portions of one orbital and a portion of a second orbital occupy the same region of space.
Fig #4, from my notes
The energy of the system depends on how much the orbitals overlap. Figure 1 below illustrates how the sum of the energies of two hydrogen atoms (the colored curve) changes as they approach each other. When the atoms are far apart, there is no overlap between their orbitals. By convention, we set the sum of the energies at zero in this state. As the atoms move together, their orbitals begin to overlap, allowing each electron to feel the attraction of the nucleus in the other atom. At the same time, the electrons and nuclei begin to repel each other.
Fig#4 interactions descriptions
a) The interaction of two hydrogen atoms changes as a function of distance. (b) The energy of the system changes as the atoms interact. The lowest (most stable) energy occurs at a distance of 74 pm, which is the bond length observed for the H2 molecule.
Fig#5, in addition to the distance of 2 orbitals, what else does orientation of orbitals affect?
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. Figure 3 below illustrates this for two p orbitals from different atoms; the overlap is greater when the orbitals overlap end to end rather than at an angle.
Fig 5 still, which orbital overlap is better and when?
The overlap of two p orbitals is greatest when the orbitals are directed end to end. (b) Any other arrangement results in less overlap. The dots indicate the locations of the nuclei.
Fig#6, What is sigma bond ?
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), as illustrated in fig 6, respectively. A σ bond is a covalent bond where the electron density concentrates along the internuclear axis; in other words, a line between the nuclei passes through the center of the overlap region. In valence bond theory, single bonds in Lewis structures are described as σ bonds.
A pi bond what is it? Fig#7
A pi bond (π bond) is a type of covalent bond that results from the side-by-side overlap of two p orbitals, as illustrated in Figure 7. 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.
Pi (π) bonds form from the side-by-side overlap of two p orbitals. The dots indicate the location of the nuclei.
Fig#8, all single bonds are ____ bonds, multiple ones consist of ______ __ and ___ bonds.
While all single bonds are σ bonds, multiple bonds consist of both σ and π bonds. As the Lewis structures below suggest, O2 contains a double bond, and N2 contains a triple bond. A double bond consists of one σ bond and one π bond. In contrast, a triple bond consists of one σ bond and two π bonds. Between any two atoms, the first bond that forms will always be a σ bond.
However, there can only be one σ bond in any given location. In any multiple bond, there will be one σ bond, while the remaining bonds will be π bonds. This chapter provides a more detailed description of these bonds later on.
Fig #8 still, average c-c single bonds is ____ kj?mo, while carbon=carbon double bond is it is ____kj/mol because….?
As seen in the table above, 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.
How are sigma and Pi bonds Similar and different?
Solution
Similarities: Both types of bonds result from overlap of atomic orbitals on adjacent atoms and contain a maximum of two electrons.
Differences: σ bonds are stronger and result from end-to-end overlap and all single bonds are σ bonds; π bonds between the same two atoms are weaker because they result from side-by-side overlap, and multiple bonds contain one or more π bonds (in addition to a σ bond).
Explain why bonds occur at specific average bond distances instead of the atoms approaching each other infinitely close.
Solution
The specific average bond distance is the distance with the lowest energy. At distances less than the bond distance, the positive charges on the two nuclei repel each other, and the overall energy increases.
Use valence bond theory to explain the bonding in F2, HF, and ClBr. Sketch the overlap of the atomic orbitals involved in the bonds.
