MORISSON   Organic Chemistry

MORISSON Organic Chemistry


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most (although int all) molecules. Furthermore, this idea closely
Fnlhls the chemist's classical concept of a bond as a force acting between two
rm and pretty much independent of the rest of the molecule; it can hardly be
cilgotal that this concept has worked amazingly well for a hundred years.
SCnificantly, the exceptional molecules for which classical formulas do not work
-.l"tt thoie for which the localized molecular orbital approach does not work
ertcr. (Eventhesecases, we shallfind, canbe handledbyarathersimple adaptation
dclrssicalformulas, anadaptationwhichagainparallelsamethodof mathematical
groximation.)
Tbe second assumption, of a relationship between'atomic and molecular
Ititals, is a highly reasonable one, as discussed in the following section. It has
Fvco so useful that, when necessary, atomic orbitals of certain kinds have been
rurrad just so that the assumption can be retained.
l, Thecovalentbond
Now let us consider the formation of a molecule. For conVenience we shall
inse this as happening by the coming together of the individual atoms, altholgh
Ln molecules aie not actually made this way. We make physical models of
olccules out of wooden or plastic balls that represent the various atoms; the
l\u20action of holes or snap fasteners tells us how to put them together. In the same
rry. we shall make mental models of molecules out of mental atoms; the locatio'n
of iomic orbitals-some of them imaginary-will tell us how to put these together.
For a covalent bond to form, two atoms must be located so that an orbital of
* mlaps an orbital of the other; each orbital must contain a single electron.
l'bbbbbbbbbbbbbbbb\ufffd:o this happens, the two atomic orbitals merge to form a single bond orbital
rlcf il ocpufieO by both electrons. The two electrons that occupy a bond orbital
Ir beve opposite spins, that is, must be paired. Each electron has available to it
-.rire mnd orbital, and thus may be considered to 
"belong to" both atomic
t|irerrangement of electrons and nuclei contains less energJ-that is, is mot\u20ac
)-tD the arrangement in the isolated atoms; as a result, formation of a
npanied by evolution of engrgy' The amount of energy (permole) that
rhcn a bond is formed (orthe amount that must be put into bre,ak the
tfu bond dissociation ercrgy. Fot a given pair of atoms, 6e grcater
rlomic orbitals, the stronger the bond.
lo STRUCTURE AND PROPERTIES CHAP. T
What gives the covalent bond its strength? It is the increase in electrostatic
attraction. In the isolated atoms, each electron is attracted by-and attracts--onepcitive nucleus; in the molecule, each electron is attracted by t*opositive nuclei.
It is the concept of " overlap " that provides the mental UiiOge between atomic
oditals and bond orbitals. Overlap of atomic orbitals means tha:t trre bond orbital
occupies much of the same region in space that was occupied by both atomic
orbitals. consequently, an electron from one atom can, to aionsiderable extent,
remain in its original, favorable location with respect to .. its,' nucleus, and at the
sane time occupy a similarly favorable location with respect to the second nu"teur;
the same holds, of course, for the other electron.
- .The princ ipleof maximumoaerlap,firststated in l93l by Linus pauling (at the California
Pit^tll1gfl1$notoev), has.been ianked-only slightly 6d.*iil;;;i'"iion principle inrmportance to the understanding of molecular structure.
As our fust example, let us consider the formation of the hydrogen molecule,
H2, from-two hydrogen atoms. Each hydrogen atom has one electron, which
occupies the ls orbital. As we have seen, this is orbital is a sphere with its center
at the atomic nucleus. For a bsnd to form, the two nuclei must be brought closely
enough together f,or overlap of the atomic orbitals to occur (Fig. I . 3). For hydrogen,
the system is most stable when the distance betweenthe iuclei is},\iA;iti.
distance is called the bond length. At this distance the stabilizing effect of overlap
is exactly balanced by repulsion between the similarly charged nu-"I"i. Th. ,.sultini
hydrogen molecule contains 104 kcal/mol less energy thin the hydrogen atomifrom which it was made. y: ryv that the hydrogen-irydrogen bond has a lengthof 0.74A and a strength of 104 kcal. 
J
ligure 1.3 Bgnd formation: H2 molecule. (a) Separate s orbitals.(D) Overlap ofs orbitais. (c) and-(d) The o bond orUitut.
This bond orbital has roughly the shape we would expect from the merging of
two s orbitals. As shown in Fig. 1.3, it is sausage-shaped, with its long "*i, tyirrg
alqng the line joining the nuclei. It is cylindrically symmetical about this ionl
aris; that is, a slice of the sausage is circular. nond orlbitals having this shape are
alldo uhitals (sigrna orbitals) and the bonds are called o bonds.We may visualize
thc hydrogcNt nolecule as two nuclei embedded in a single sausage-shapid electron
clqrd. Thiscloudisdensest inthe region between the twlnuclei, ivhereihe negative
charge is anracted moat strongly by the two positive charges.
- lF sire of the hydrogen molecule-as measured, ray, by the volume insidethe 951 probability surface-is considerably smaller thanthat bf a single hydrogen
r t t IIYBRID OI'BITAIS: qp l l
t-- Ahhough surprising at,first, this shrinking of the electron cloud is actually
fi ruuld be expected. It is'the powerful attraction of the electrons by twonuclei
L gives the molecule greater stability than the isolated hydrogen atoms; this
Dcan that the electrons are held tighter, closer, than in the atoms.
; Next, let us consider,the formation of the fluorine molecule, F2, from two
-dDc atoms. As we can see from our table of electronic configurations (Table
LIL a fluorine atom has two electrons in the ls orbital, two electrons in the fu
Iti3d, and two electrons in each of two 2p orbitals. In the third 2p orbitalthere is
r.ingle electron which is unpaired and available for bond formation. overlap of
Cirp orbital with a similarp orbital of another fluorine atom permits electrons to
pir and the bond to form (Fig. l. ). The electronic charge is concentrated between
ft two nuclei, so that the back lobe of each of the overlapping orbitals shrinks to
o'o o.o
o.o.o
(a)
(c)(b)
Figure 1.4 Bond formation : F2 molecule. (a) separatep orbitals. (D) overlap
ofp orbitals. (c) The o bond orbital.
a comparatively small size. Although formed by overlap of atomic orbitals of a
different kind, the fluorine-fluorine bond has the same general shape as the
bydrogen-hydrogen bond, being cylindrically symmetrical about a line joining the
rrrclei; it, too, is given the designation of o bond. The fluorine-fluorine bond has
e hngth of |.42 A and a strength of about 38 kcal.
As the examples show, a covalent bond results from the overlap of two atomic-
orbitals to form a bond orbital occupied by a pair of electrons. Each kind of coaalmt
btd has a characteristii length and strength.
1.9 Hybrid orbitals: sp
Let us next consider beryllium chloride, BeClr.
Beryllium (Table 1.1) has no unpaired electrons. How are we to account for
its combining with two chlorine atoms? Bond formation is an energy-releasing
o o o o
(gbilizing) process, and tile ten{ency is tdform uonds-and as many as possible-
cra if this results in bond orbitals that bear little resemblance to the atomic
atftds we have talked about. If our method of mental molecule-building is to be
Tplicd here, it must be modified. we must invent an imaginary kind of beryllium
-m, one that is about to become bonded to two chlorine atoms.
2p2sl s
o
t2 STRUCTURE AND PROPERTIES
2s 2p
/-----^-
o o c o
CHAP. I
One electron promoted:
two unpaired electrons
To arrive at this divalent beryllium atom, let us do a little electronic book-
keeping. First, we "promoti'" one of the-2s electrons to an emptyp orbital:
This provides twounpaired electrons, which are needed forbondingtotwo chlorine