Valence Bond Theory (10.2.1) Flashcards
• When two atoms are in proximity, they will arrange themselves at the lowest
possible energy level.
• When two atoms are in proximity, they will arrange themselves at the lowest
possible energy level.
• Sigma bonds are formed by the end-on overlap of orbitals along the internuclear
axis.
• Sigma bonds are formed by the end-on overlap of orbitals along the internuclear
axis.
• Pi bonds are formed by side-on overlap of orbitals on either side of the
internuclear axis.
• Pi bonds are formed by side-on overlap of orbitals on either side of the
internuclear axis.
As the distance between two hydrogen atoms
decreases, the attraction between them increases.
The reason is that the nucleus of one atom can
attract the electrons of the other atom. Beyond a
critical point, however, the nuclei will repel each
other. The diagram on the left depicts the optimal
point of attraction between the atoms, a distance of
74 picometers. At this distance, the energy of the
electrons is lower than at any other distance. Atoms
always seek the lowest possible energy level.
Valence bond theory states that a chemical bond
results from the overlap of two orbitals, each of
which is half filled. With respect to dihydrogen, the
electron from each orbital occupies the other
hydrogen’s orbital because two electrons represent
the most stable configuration, forming what is
known as a sigma bond. In contrast, helium doesn’t
form a bond because there is not a place for
another electron. The 1s orbital in helium contains
two electrons and consequently cannot accept the
electrons of another atom.
Sigma bonds form when two s orbitals overlap, as in
dihydrogen, or when an orbital that has directional
character, such as a p orbital, overlaps another
orbital along the direction of the orbital axis. For
example, the pz orbital of fluorine overlaps a 1s
orbital along its axis, resulting in a sigma bond.
As the distance between two hydrogen atoms
decreases, the attraction between them increases.
The reason is that the nucleus of one atom can
attract the electrons of the other atom. Beyond a
critical point, however, the nuclei will repel each
other. The diagram on the left depicts the optimal
point of attraction between the atoms, a distance of
74 picometers. At this distance, the energy of the
electrons is lower than at any other distance. Atoms
always seek the lowest possible energy level.
Valence bond theory states that a chemical bond
results from the overlap of two orbitals, each of
which is half filled. With respect to dihydrogen, the
electron from each orbital occupies the other
hydrogen’s orbital because two electrons represent
the most stable configuration, forming what is
known as a sigma bond. In contrast, helium doesn’t
form a bond because there is not a place for
another electron. The 1s orbital in helium contains
two electrons and consequently cannot accept the
electrons of another atom.
Sigma bonds form when two s orbitals overlap, as in
dihydrogen, or when an orbital that has directional
character, such as a p orbital, overlaps another
orbital along the direction of the orbital axis. For
example, the pz orbital of fluorine overlaps a 1s
orbital along its axis, resulting in a sigma bond.
Pi bonds form by the sideways overlap of two
parallel p orbitals. Pi bonds are not as strong as
sigma bonds.
Nitrogen has an electron configuration of 1s2
2s2
2p3
.
Because there are three unpaired electrons in
nitrogen, three bonds can be made. There is one
sigma bond and two pi bonds. There is side-on
overlap along both the x and y axes (the two pi
bonds) and end-on overlap along the z-axis (the
sigma bond). This bond is a triple bond, as
predicted by Lewis dot structure.
Pi bonds form by the sideways overlap of two
parallel p orbitals. Pi bonds are not as strong as
sigma bonds.
Nitrogen has an electron configuration of 1s2
2s2
2p3
.
Because there are three unpaired electrons in
nitrogen, three bonds can be made. There is one
sigma bond and two pi bonds. There is side-on
overlap along both the x and y axes (the two pi
bonds) and end-on overlap along the z-axis (the
sigma bond). This bond is a triple bond, as
predicted by Lewis dot structure.
Look at the graph that compares the internuclear separation between two hydrogen atoms with their potential energy. What is indicated by Point 1 on the graph?
This is the point at which the two hydrogen atoms have the lowest potential energy and a defined internuclear axis length. (C)
At this point, the two hydrogen atoms have assumed an ideal distance for sharing their lone electrons, so their potential energy is minimized and there is a specific internuclear axis length value. If that value increases or decreases, the potential energy rapidly increases.
he bonding behavior of two hydrogen atoms can be observed in the graph.
Which of the following statements is completely correct?
Electron sharing actually can reduce the potential energy of the system, which in this case is H2. The potential energy is minimized at an ideal internuclear separation distance of 74 pm. At any other internuclear distance you do not have minimal potential energy for the system. (C)
There is only one point on the curve (at 74 pm) where there is a minimum potential energy value. If you move in either direction along the curve (i.e., increasing or decreasing the internuclear separation distance), you increase the potential energy of the system.
Oxygen has an electron configuration of 1s2 2s2 2p4. Which statement about oxygen is not correct?
An oxygen atom would be less likely to form a bond than an atom that has an electron configuration of 1s2 2s2 2p6. (D)
An atom that has an electron configuration of 1s2 2s2 2p6 has three filled electron orbitals, so it will be much less reactive than oxygen, which has two unpaired electrons in its outer p shell. This is because its electron configuration is 1s2 2s2 2p4. The outer p shell has two sites for pair electrons. Since there are only four (of six possible) electrons in that orbital, there is only one pair of electrons in that orbital. So oxygen will tend to want to interact and form bonds with other atoms that have lone electron pairs.
Which statement about the bonds in N2 is not correct? Remember that the electron configuration for nitrogen in the ground state is 1s22s22p3.
There are two σ bonds and one π bond. The σ bonds are arranged along the x- and y-axes, and the π bond is arranged along the z-axis. (A)
The opposite is true; there are one σ bond and two π bonds.
Which statement about the chemical bond between two hydrogen atoms is not true?
The amount of energy needed to break the bond is twice the potential energy of the ideally bonded atoms. (C)
This statement would be correct if you eliminated the term twice. In other words, the amount of energy needed to break the bond equals the potential energy of the ideally bonded atoms.
Suppose that a theoretical gas atom has an electron configuration of 1s22s22p6. Which statement best describes this atom’s electron arrangement and theoretical reactivity with an atom that has an unpaired electron?
This atom has five orbitals, 1s, 2s, and three 2p orbitals. They each contain paired electrons. This atom will not display highly reactive tendencies, because its orbitals are already filled. (B)
This atom has a stable electron arrangement because it has two sets of paired electrons in its s orbitals. This is actually the electron configuration for neon.
What is the difference between a σ bond and a π bond?
A σ bond’s overlap occurs along the internuclear axis. (B)
This is what distinguishes a σ bond from a π bond (which has two overlap sites at areas above and below the line defined by the two atomic nuclei).
Which statement best describes the bonding in an O2 molecule?
O2 has one σ bond and one π bond. (C)
Because oxygen has an electron configuration of 1s22s22p4, oxygen accepts two electrons to fill its valence shell. Two oxygen atoms can fill their valence shells by forming a σ bond and a π bond.
Which statement about valence bond theory is not true?
An atom with a filled valence shell is unstable and therefore very willing to form a covalent bond with another atom. (D)
This statement is not true. The atom does not have an unpaired electron available, so it will be highly unreactive (and therefore stable).
The electron configuration for fluorine, F, is 1s2 2s2 2p5 . Which statement about the electron configuration for fluorine is not correct?
All the orbitals in the fluorine atom have complete electron pairing. (A)
The z-axis site in the 2p orbital has only one assigned electron, not an electron pair. You know this occurs on the z-axis because of the convention of assigning electrons to p sites in axis order of x, y, then z.
