BONDS AND BONDING IN ATOMS AND
MOLECULES
A HISTORY OF IDEAS
BEGINNINGS
Democritus: Universe was atoms and empty space
Atoms –indestructible; indivisible; in constant motion; connected by
hooks and eyes
Decline of Roman Empire – decline in atomism
Earth, air, fire and water; alchemy and effort to change base metal to gold
Atomism arose again in the 17th C.
Boyle - matter consisted of corpuscules that arrange
themselves in groups or clusters to provide the substances
familiar in life.
Newton in his Principia of 1704 - particles of matter attract
one another by some force which in immediate contact is
very strong and at small distances causes the
chemical operations.
Developments
Ettiene Geoffray and affinity tables 1718
1803 John Dalton defines atomic weight - every
element was referred to the lightest element, hydrogen
Berzelius – introduced modern symbols for elements;
Compounds composed of atoms held together by alternating +ve
and –ve charges
Edward Frankland (1825 – 1891)
1852 - proposed that atoms of each element have a
definite ‘saturation capacity’
i.e. the can only combine with a limited number of
other atoms
forerunner of ‘valency’
Structural Chemistry
Origins
Kekule
1857 - Valency; the combining power of an element, as
measured by the number of hydrogen atoms it can
displace or combine with - proposed fixed valencies
for elements. Notably carbon has a valency of 4
Couper
Pasteur (1822 – 1895)
Tartaric acid from dregs of wine rotated plane of polarization laboratory produced tartaric acid no effect
1874 van’t Hoff - optical activity was due to the bonds in
carbon atoms being directed towards the corners of a
tetrahedron.
D-Tartaric Acid
L-Tartaric Acid
August von Hoffman
First Electronic Ideas
1897 discovery of the electron by J J Thomson
Gilbert Newton Lewis (1875 – 1946)
1902 cubical model of the electronic
structure of atoms
Lewis noted that elements with certain
number of electrons had stability
– usually 8 electrons.
when a layer of 8 electrons was complete a new
layer is started
placed the electrons at the corners of a cube.
A single bond between two atoms was formed
when two cubes shared an edge; a double bond
when they shared a face.
key concept - sharing of electrons between atoms to form the covalent
bond.
Ionic bond was formed when the two shared electrons became localised on
one atom e.g. NaCl – Na+Cl-
Irving Langmuir - developed mathematical
equations to predict the number of electron pair
bonds
After 1921-22 no more contributions
1923, Lewis published ‘Valence and the Structure of Atoms and
Molecules’ and then left the field
The major distillations from Lewis’s ideas are;
Atoms in compounds tend to assume the electron configuration of noble gases through
sharing of electrons or electron transfer.
Electrons may form part of a shell of two different atoms and cannot be said to belong to
one or the other exclusively (but Lewis viewed the shared electrons in a bond as being fixed
between the atoms).
The shell concept needed Bohr in the 1920s to square his original model with the new QM
descriptions of electron distribution in atoms
Energy levels and shells
Electrons are arranged in different shells around the nucleus. Each succeeding shell can only hold a certain
number of electrons before it becomes full..
Maximum capacity of the first three shells
energy level or
shell
maximum number of
electrons
first
2
second
8
third
8
lithium atom has three electrons. Two are in the first energy level, and one in the second.
carbon atom has six electrons. Two are in the first energy level, and four in the second energy level.
Arrangement of electrons in a lithium atom
Arrangement of electrons in a carbon atom
.
Valence Bond Theory
Schrodinger’s wave theory of the energy and distribution of electrons in atoms, three quantum
numbers arise naturally;
n – principal quantum number – defines the energy of the electron shell
l – orbital angular momentum number -values 0 to n-1 defines the shape of the orbital
ml – magnetic quantum number – values -l to +l defines the orientation in space of the orbital
Each orbital can contains a maximum of two electrons
Pauli exclusion principle –no two electrons can have the same quantum
numbers
fourth quantum number– spin quantum number ms with values of -1/2 and
+1/2
n
l
ml
Max
electrons
ms
0.5
1
2pz orbital
2 electrons
2py orbital
2 electrons
2px orbital
2 electrons
2s orbital
2 electrons
1s orbital
2 electrons
-0.5
0.5
1
0
-0.5
0.5
2
-1
-0.5
0.5
0
0
-0.5
0.5
1
0
0
-0.5
Covalent bond in H2 - two 1s orbitals overlap - electron density distribution
greater between the two nuclei
sigma (σ) bond
Fritz London
Linus Pauling (1901-1994)
chemical bond formation was the sharing of
electrons between atoms
Walter Heitler
effort of the bonded atoms to have a
completed outer shell of 8 electrons
types of bonds formed encountered in the nitrogen molecule N2;
nitrogen atom has the valence shell electronic structure 2s(2); 2px(1); 2py(1); 2pz(1).
Sigma bond is formed by the overlap of the 2pz orbitals on each nitrogen atom.
The remaining 2p orbitals are aligned with each other and form pi bonds by spin pairing of the
electrons in each orbital.
Pauling introduced two major concepts into valence bond theory;
Resonance –molecules can have structures of identical energy the
wave functions of each are superposed. A good example is the
resonance structures of benzene
Hybridisation – the valence shell configuration of carbon is 2s(2); 2px(1); 2py(1)
Pauling suggested one of the 2s electrons was promoted to the 2pz orbital
then all the orbitals mixed together to form 4 hybrid orbitals
– 1 part s and 3 parts p orbitals hence sp3 orbitals
each of these orbitals points in the direction of a regular tetrahedron
NP in Chemistry in 1954 – "for his research into the
nature of the chemical bond and its application to the
elucidation of the structure of complex substances"
Molecular Orbital Theory
John Lennard-Jones (1894 – 1954). Friedrich Hund (1896 – 1997);
Robert Mulliken (1896 – 1986);
Basis of MO theory - linear combination of atomic orbitals
Major difference with VB theory - electrons do not belong to
particular bonds - spread throughout the molecule
Metals
strength of metallic bonds varies
dramatically.
caesium melts at 28.4C, mercury is a
liquid at room temperature, but tungsten
melts at 3680C.
Metallic bonds tend to be weakest for
elements that have nearly empty (as in
Cs) or nearly full (Hg) valence subshells,
and strongest for elements with
approximately half-filled valence shells
(as in W)
band structure also accounts for metal lustre
Silver reflects light with equal efficiency for all wavelengths –
many unoccupied molecular orbitals populated by light of all wavelengths same frequency when the electrons drop back to lower levels.
Copper - fewer unoccupied energy levels populatable by violet,
blue or green light absorption - light emitted when electrons drop back is at
lower frequencies - yellow, orange and red
Hydrogen Bond
occurs when a hydrogen (H) atom covalently bound to a highly electronegative atom such as
nitrogen (N), oxygen (O), or fluorine (F) experiences the electrostatic field of another highly
electronegative atom nearby
electronegative atom attracts the electron cloud from around the hydrogen leaving the atom with a
positive partial charge.
Small size of hydrogen means the resulting charge provides a large charge density.
This strong positive charge density attracts a lone pair of electrons on another heteroatom, forming a
hydrogen bond
•
F−H···:F (161.5 kJ/mol)
•
O−H···:N (29 kJ/mol)
•
O−H···:O (21 kJ/mol)
•
N−H···:N (13 kJ/mol)
•
N−H···:O (8 kJ/mol)
Hydrogen bonding phenomena
•
Dramatically higher boiling points of NH3, H2O, and HF compared to the heavier
analogues PH3, H2S, and HCl.
•
High water solubility of many compounds such as ammonia - hydrogen bonding with
water molecules.
•
Ice is less dense than liquid water - due to a crystal structure stabilized by hydrogen
bonds.
•
Strength of nylon and cellulose fibres.
•
Wool, a protein fibre, held together by hydrogen bonds, providing recoil when stretched.
Washing at high temperatures breaks the hydrogen bonds leading to permanent loss of shape
.
Bond Energies /kJ/mol
Single Bonds
H—H
H—F
H—Cl
H—Br
H—I
432
565
427
363
295
C—H
C—C
C—N
C—O
C—F
C—Cl
C—Br
C—I
C—S
413
347
305
358
485
339
276
240
259
N—H
N—N
N—F
N—Cl
N—Br
N—O
O—H
O—O
O—F
O—Cl
O—I
391
160
272
200
243
201
467
146
190
203
234
F—F
F—Cl
F—Br
Cl—Cl
Cl—Br
154
253
237
239
218
Multiple Bonds
I—I
I—Cl
I—Br
149
208
175
S—H
S—F
S—Cl
S—Br
S—S
347
327
253
218
266
Si—Si
Si—H
Si—C
Si—O
340
393
360
452
C=C
C≡C
O=O
C≡O
N=O
N=N
N≡N
C≡N
C=N
614
839
495
1072
607
418
941
891
615
SOURCES
Clayden, Greeves et al ‘Organic Chemistry’
Atkins and de Paula ‘Physical Chemistry’
Partington ‘A Short History of Chemistry’
Coffey ‘Cathedrals of Science’
http://scarc.library.oregonstate.edu/digitalresources/pauling/
http://www.quantum-chemistry-history.com/Ueberb1.htm
http://academictree.org/chemistry/index.php