Chapter 1: Structure and Bonding
Atomic Structure
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Structure & Bonding |
The study of organic chemistry must at some point extend to the molecular level, for the physical and chemical properties of a substance are ultimately explained in terms of the structure and bonding of molecules. This module introduces some basic facts and principles that are needed for a discussion of organic molecules.
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Electron Configurations |
Electron Configurations in the Periodic Table
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Four elements, hydrogen, carbon, oxygen and nitrogen, are the major
components of most organic compounds. Consequently, our understanding of
organic chemistry must have, as a foundation, an appreciation of the electronic
structure and properties of these elements. The truncated periodic table shown
above provides the orbital electronic structure for the first eighteen elements
(hydrogen through argon). According to the Aufbau principle, the
electrons of an atom occupy quantum levels or orbitals starting from the lowest
energy level, and proceeding to the highest, with each orbital holding a
maximum of two paired electrons (opposite spins).
Electron shell #1 has the lowest energy and its s-orbital is the first to be filled. Shell #2 has four higher energy orbitals, the 2s-orbital being lower in energy than the three 2p-orbitals. (x, y & z). As we progress from lithium (atomic number=3) to neon (atomic number=10) across the second row or period of the table, all these atoms start with a filled 1s-orbital, and the 2s-orbital is occupied with an electron pair before the 2p-orbitals are filled. In the third period of the table, the atoms all have a neon-like core of 10 electrons, and shell #3 is occupied progressively with eight electrons, starting with the 3s-orbital. The highest occupied electron shell is called the valence shell, and the electrons occupying this shell are called valence electrons.
The chemical properties of the elements reflect their electron configurations. For example, helium, neon and argon are exceptionally stable and unreactive monatomic gases. Helium is unique since its valence shell consists of a single s-orbital. The other members of group 8 have a characteristic valence shell electron octet (ns2 + npx2 + npy2 + npz2). This group of inert (or noble) gases also includes krypton (Kr: 4s2, 4p6), xenon (Xe: 5s2, 5p6) and radon (Rn: 6s2, 6p6). In the periodic table above these elements are colored beige.
The halogens (F, Cl, Br etc.) are one electron short of a valence
shell octet, and are among the most reactive of the elements (they are colored
red in this periodic table). In their chemical reactions halogen atoms achieve
a valence shell octet by capturing or borrowing the eighth electron from
another atom or molecule. The alkali metals Li, Na, K etc. (colored
violet above) are also exceptionally reactive, but for the opposite reason.
These atoms have only one electron in the valence shell, and on losing this
electron arrive at the lower shell valence octet. As a consequence of this electron
loss, these elements are commonly encountered as cations (positively charged
atoms).
The elements in groups 2 through 7 all exhibit characteristic reactivities and
bonding patterns that can in large part be rationalized by their electron
configurations. It should be noted that hydrogen is unique. Its location in the
periodic table should not suggest a kinship to the chemistry of the alkali
metals, and its role in the structure and properties of organic compounds is
unlike that of any other element.
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Bonding & Valence |
Chemical Bonding and Valence
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2Na + Cl2 |
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2NaCl |
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2H2 + O2 |
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2H2O |
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C + O2 |
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CO2 |
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C + 2F2 |
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CF4 |
As noted earlier, the inert gas elements of group 8 exist as monatomic gases, and do not in general react with other elements. In contrast, other gaseous elements exist as diatomic molecules (H2, N2, O2, F2 & Cl2), and all but nitrogen are quite reactive. Some dramatic examples of this reactivity are shown in the following equations.
Why do the atoms of many elements interact with each other and with other elements to give stable molecules? In addressing this question it is instructive to begin with a very simple model for the attraction or bonding of atoms to each other, and then progress to more sophisticated explanations.
Ionic Bonding
When
sodium is burned in a chlorine atmosphere, it produces the compound sodium chloride.
This has a high melting point (800 °C) and dissolves in water to give a
conducting solution. Sodium chloride is an ionic compound, and the crystalline
solid has the structure shown on the right. The transfer of the 2s electron of
a sodium atom to the half-filled 2p orbital of a chlorine atom generates a
sodium cation (neon valence shell) and a chloride anion (argon valence shell).
Electrostatic attraction results in these oppositely charged ions packing
together in a lattice. The attractive forces holding the ions in place can be
referred to as ionic bonds.
Covalent
Bonding
The other three reactions shown above give products that are very different
from sodium chloride. Water is a liquid at room temperature; carbon dioxide and
carbon tetrafluoride are gases. None of these compounds is composed of ions. A
different attractive interaction between atoms, called covalent bonding, is
involved here. Covalent bonding occurs by a sharing of valence electrons,
rather than an outright electron transfer. Similarities in physical properties
(they are all gases) suggest that the diatomic elements H2, N2,
O2, F2 & Cl2 also have covalent bonds.
Examples of covalent bonding shown below include hydrogen, fluorine, carbon
dioxide and carbon tetrafluoride. These illustrations use a simple Bohr
notation, with
valence electrons designated by colored dots. Note that in the
first case both hydrogen atoms achieve a helium-like pair of 1s-electrons by
sharing. In the other examples carbon, oxygen and fluorine achieve neon-like
valence octets by a similar sharing of electron pairs. Carbon dioxide is
notable because it is a case in which two pairs of electrons (four in all) are
shared by the same two atoms. This is an example of a double covalent bond.
These electron sharing diagrams are a useful first step in understanding covalent bonding, but it is quicker and easier to draw Kekulé formulas in which each shared electron pair is represented by a line between the atom symbols. Non-bonding valence electrons are shown as dots. Some examples of these structural formulas are given in the following table.
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Multiple bonding, the sharing of two or more electron pairs, is illustrated by ethylene and formaldehyde (each has a double bond), and acetylene and hydrogen cyanide (each with a triple bond). Boron compounds such as BH3 and BF3 are exceptional in that conventional covalent bonding does not expand the valence shell occupancy of boron to an octet. Consequently, these compounds have an affinity for electrons, and they exhibit exceptional reactivity when compared with the compounds shown above.
Valence
The number of valence shell electrons an atom must gain or lose to achieve a
valence octet is called valence. In covalent compounds the number of bonds
which are characteristically formed by a given atom is equal to that atoms
valence. From the formulas written above, we arrive at the following general
valence assignments:
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Atom |
H |
C |
N |
O |
F |
Cl |
Br |
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Valence |
1 |
4 |
3 |
2 |
1 |
1 |
1 |
VBT (Valence Bond Theory)
Valence bond theory
is a model that can be used to determine the location of the electrons within
the central atom. By spreading out the electrons within the atom, bonds can be
formed with the attached atoms. The process
of spreading out electrons creates new orbital types called hybrid orbitals.
There are several types of hybrid orbitals that can be formed from s, p, and d
type atomic orbitals.
Atomic
Orbitals Hybrid Molecular
Mixed orbital Shape
s and p sp Linear
s, p and p sp2 trigonal
planar
s,p,p and p sp3 tetrahedral.
d,s,p,p and p dsp3 trigonal bipyramidal
d,d,s,p,p and p d2sp3 octahedral
We can use box
diagrams to indicate where the electrons are within the central atom. Consider
carbon.
s p
When carbon is
hybridized to form bonds with say, 4 hydrogens, the orbital set looks like:
H H
H H
sp3 sp3
Write the box diagrams for ClF3.
Cl=
s p d
F
F F
Cl=
dsp3 d
Write the box diagram
for XeF4.
Xe=
s p d
F F
F F
Xe=
d2sp3 d
DOUBLE BONDING AND VBT
VBT makes the
explanation of multiple bonds rather simple.
Consider the compound CH2O. The Lewis structure reveals a
double bond in the structure. We can
see that the compound is going to be trigonal planar.
The Box diagram for
carbon is:
C
s p
In order for
carbon to bond with three elements, hybridization must occur, but only 3
bonding orbitals are needed. This yields the box diagram:
H
H :O: p
bond
C
sp2 p
with 3 sp2
orbitals and a left over p-orbital.
Consider the oxygen
in the atom in H2CO.
O
s p
Oxygen hybridizes and
forms a single sigma (s) bond with
the oxygen. Oxygen, like carbon, also has a left over p-orbital to form a pi (p) bond.
.C. p
bond
O
sp2 p
Both oxygen and carbon
contain an extra electron in their unhybridized p-orbitals. Since both carbon
and oxygen are small, these p-orbitals have effective overlap and a p bond is formed.
MOT
Molecular orbital theory
attempts to describe the bonding in electrons with the concept of molecular
orbitals. Molecular orbitals consist of
linear combinations of atomic orbitals, much like the hybrid orbitals, but the
molecular orbitals are derived for all the outer orbitals of all the atoms in
the compound. Molecular orbitals are similar to atomic orbitals. For example,
two electrons can occupy any molecular orbital, the electrons must have
opposite spin and molecular orbitals have specific energies. However, when two atomic orbitals combine to
form a molecular orbital, two molecular orbitals are formed. One of the
molecular orbitals is a bonding orbital and when occupied represents a lowering
of the overall energy of the electrons.
The other molecular orbital is an antibonding orbital, and when occupied
represents an increase in the energy within the molecule. Consider the formation of H2 from
two hydrogen atoms.
s*1s
1s
1s
H H
s 1s
H2
The formation of H2
then represents a lowering of the overall energy of the molecule.
Consider the
formation of He2:
s*1s
1s
1s
H H
s
1s
He2
We can see that the
energy reduction of electrons in the bonding molecular orbitals is offset by
the gain in energy of the electrons in the antibonding orbitals. There is no reduction in energy for the
formation of He2 and therefore no reason for it to form. Therefore He2 does not exist.
The bond order is the
total number electron in bonding orbitals minus the total number of electrons
in antibonding orbitals, the whole thing divided by 2.
Consider the
formation of bonds in N2.
Atomic orbitals of
N Molecular orbitals of N2 Atomic orbitals of N
s*2p
p*2p
2p 2p
s2P
p2P
s*2s
2s 2s
s2s
s*1s
1s 1s
s1s
From the energy
diagram we can see that the bond order for N2 is 3, which is
confirmed by Lewis structures. Draw
this Lewis structure.
What is the molecular
orbital diagram for O2?
Atomic orbitals of
O Molecular orbitals of O2 Atomic orbitals of O
s*2p
p*2p
2p 2p
s2P
p2P
s*2s
2s 2s
s2s
s*1s
1s 1s
s1s
What is the bond order for O2?
What is the molecular
orbital energy diagram for F2? Determine the bond order.
Atomic orbitals of
F Molecular orbitals of F2 Atomic orbitals of F
s*2p
p*2p
2p 2p
s2P
p2P
s*2s
2s 2s
s2s
s*1s
1s 1s
s1s
Examine the molecular
orbital diagrams for O2 and F2. We can see that O2
has two unpaired electrons while F2 contains no unpaired
electrons. When compounds have unpaired
electrons they have the magnetic property called paramagnetism. Paramagnetic
substances are attracted into a magnetic field. When all the electrons in a compound are paired, the compound is
said to be diamagnetic. Diamagnetic substances are weakly pushed out of a
magnetic field.
The Lewis structure
for O2 fails to predict the paramagnetic properties of the compound,
but MOT describes not only the bond order, but also the paramagnetic
properties.
Hybridization: sp3 Orbitals and the
Structure of Methane
The Structure of Ethane
Hybridization: sp 2 Orbitals and the
Structure of Ethylene
Hybridization: sp Orbitals and the Structure of
Acetylene
Hybridization of Other Atoms: Nitrogen and Oxygen