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 Couper

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

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-

1923, Lewis published ‘Valence and the Structure of Atoms and

Molecules’ and then left the field

Irving Langmuir - developed mathematical

equations to predict the number of electron pair

bonds

After 1921-22 no more contributions

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 shellsElectrons 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

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

.

energy level or

shell

maximum number of

electrons

first 2

second 8

third 8

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 ms Max

electrons

0.5

1

2pz orbital 2 electrons

-0.5

0.5

1 0

2py orbital 2 electrons

-0.5

0.5

2

-1

2px orbital 2 electrons

-0.5

0.5

0 0

2s orbital 2 electrons

-0.5

0.5

1 0 0

1s orbital 2 electrons

-0.5

Fritz London

Walter Heitler

Covalent bond in H2 - two 1s orbitals overlap - electron density distribution

greater between the two nuclei

Linus Pauling (1901-1994)

chemical bond formation was the sharing of

electrons between atoms

effort of the bonded atoms to have a

completed outer shell of 8 electrons

sigma (σ) bond

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

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

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)

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 Multiple Bonds

H—H 432 N—H 391 I—I 149 C = C 614

H—F 565 N—N 160 I—Cl 208 C ≡ C 839

H—Cl 427 N—F 272 I—Br 175 O = O 495

H—Br 363 N—Cl 200 C ≡ O 1072

H—I 295 N—Br 243 S—H 347 N = O 607

N—O 201 S—F 327 N = N 418

C—H 413 O—H 467 S—Cl 253 N ≡ N 941

C—C 347 O—O 146 S—Br 218 C ≡ N 891

C—N 305 O—F 190 S—S 266 C = N 615

C—O 358 O—Cl 203

C—F 485 O—I 234 Si—Si 340

C—Cl 339 Si—H 393

C—Br 276 F—F 154 Si—C 360

C—I 240 F—Cl 253 Si—O 452

C—S 259 F—Br 237

Cl—Cl 239

Cl—Br 218

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.htmhttp://academictree.org/chemistry/index.php