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"Valence and Molecular Structure," Lectures 1 and 2.

"Valence and Molecular Structure," Lectures 1 and 2. 1957.
Produced for the Institutes Program of the National Science Foundation. Robert and Jane Chapin, producers.

Lecture 2, Part 5. (8:45)

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Transcript

Linus Pauling: Now, let us discuss the modern aspects of valence theory, structure theory. First, let me say that the classical structure theory, the older structure theory, has not been discarded. There has not been a revolution of such a nature that the old has been thrown out and the new has come in. Classical structure theory is still valid. There have been some improvements in structure theory, some problems of molecular structure and valence that were hard to discuss, hard to understand before, can now be discussed in a reasonable and sensible way because of the additions that have been made. Ideas about hyberdized bond orbitals, about the theory of resonance, partial ionic character of covalent bonds have come in and have made chemical structure theory more powerful.

The whole theory, classical structure theory and modern structure theory, have a sound base in experiment. It has, it has been developed largely by induction from the tens of thousands of chemical facts with, in the case of the modern development, a little help, or considerable help, I should say, from ideas that have been suggested by the theory of quantum mechanics for which we are indebted to the physicists.

The modern theory began in 1916 when Professor Lewis introduced the idea of the shared electron pair chemical bond and much contribution was made, also, by Irving Langmuir. The principle...statements that we can make about a chemical bond now are that in order to form a chemical bond between two atoms, you need to have an orbital for each atom. Let’s say atom A must have an orbital, atom B must have an orbital, and two electrons are involved, which I may write in this way, and their spins must be opposed. One orbital for each of the two atoms and a pair of electrons with opposed spins, which serve to hold the atoms together.

Here, I have a drawing representing the structure of the hydrogen molecule, H2. The two nuclei, the two protons, are at these positions, seventy-four hundredths of an angstrom apart, and the two electrons are distributed in space roughly as shown here, with a good concentration right in the region between the two nuclei. It is almost as if the nuclei were ball bearings, steel ball bearings around which some rubber has been vulcanized to hold these ball bearings firmly at this distance apart, does not permit them to escape from one another. This is the standard Lewis symbol for the hydrogen molecule, H2. The two electrons are shown between the symbols for the two hydrogen atoms.

We see that we can say that each of the hydrogen atoms has succeeded in obtaining the helium structure. This orbital for hydrogen, the 1s orbital, is occupied by the pair of electrons, which is, also occupies the 1s orbital for the other hydrogen atom.

The idea that you must have a stable orbital for each atom in order to form a bond and a pair of electrons permits a considerable addition to the power of chemical structure theory. Here, I have a drawing representing the electronic structure of the water molecule. The water molecule, H2O, has the Lewis symbol as shown here. The pair of electrons in the helium shell for oxygen is not indicated, only those in the neon shell. Here is an unshared pair of electrons occupying one orbital, a second unshared pair occupying a second orbital, a third shared pair, in this case, occupying the third orbital, and a shared pair occupying the fourth orbital. The oxygen atom now has four electron pairs, eight electrons in its neon shell. It has achieved the structure of the neon atom by sharing electrons, and the hydrogen atoms, as before, using their 1s orbitals, have achieved the helium configuration.

Many structures, many molecules, have structures such that each atom achieves the electronic structure of the nearest noble gas with two, ten, eighteen, thirty-six, fifty-four, and so on, electrons. The oxygen-hydrogen distance is known. 0.965 angstroms, the angle between the oxygen atoms is know experimentally, a hundred and four degrees, thirty minutes.

These distances are interesting in that they, the distances, as determined by spectroscopic or diffraction measurements, in that they give us an idea about the significance of the chemical bond. The carbon-carbon distance, well, let me look at this, the chlorine-chlorine distance, in the chlorine molecule, Cl2, is 1.98 angstrom. We can write the Lewis formula Cl, Cl, and I shall show also the electrons in the valence shell, the argon shell of each chlorine atom. I have written here a line like this to represent the two electrons that are shared between the chlorine atoms.

Chlorine has only sixteen electrons, that is seven in the argon shell, but by sharing a pair each chlorine atom succeeds in achieving the argon structure with eighteen electrons. The distance between the two chlorine atoms is 1.98 angstrom, as determined spectroscopically and by electron diffraction. The distance between the two carbon atoms in ethane, the carbon–carbon distance in ethane...I’ll change this to ethane, CH3, is 1.54 angstrom. 1.54 angstrom.

Now, when we examine carbon tetrachloride, by spectroscopic methods or electron diffraction, we find that the carbon-chlorine distance in carbon tetrachloride, is 1.76 angstrom. Carbon, four bonds to chlorine, I might as well continue to give the example. Well, for this chlorine atom, here again, by forming the bond, chlorine has achieved the argon structure and the carbon atom has achieved the neon structure. This distance is found to be 1.76 angstrom. 1.76 angstrom, 1.98 angstrom in chlorine, plus 1.54 angstrom in ethane, that is two, 3.52, half of that is 1.76. So, the carbon-carbon bond has a length just midway between the lengths of the carbon, the carbon-chlorine bond has a length just midway between the lengths of the carbon-carbon single bonds in ethane, and the carbon, chlorine-chlorine bonds in the chlorine molecule.

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Creator: Linus Pauling
Associated: Robert Chapin, Jane Chapin, National Science Foundation, Irving Langmuir, G. N. Lewis
Clip ID: 1957v.1-14

Full Work

Creator: National Science Foundation
Associated: Linus Pauling, Robert Chapin, Jane Chapin

Date: 1957
Genre: video
ID: 1957v.1
Copyright: More Information

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