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whiskey in the jar
Molecular structures may be a mainstay of high school science classes, but the familiar picture of balls and sticks representing atoms and the bonds among them is largely a conventional fiction. The trouble is that scientists disagree on what a more accurate representation of molecules should look like.
In the 1920s physicists Walter Heitler and Fritz London showed how to describe a chemical bond using the equations of then nascent quantum theory, and the great American chemist Linus Pauling proposed that bonds form when the electron orbitals of different atoms overlap in space. A competing theory by Robert Mulliken and Friedrich Hund suggested that bonds are the result of atomic orbitals merging into “molecular orbitals” that extend over more than one atom. Theoretical chemistry seemed about to become a branch of physics.
Nearly 100 years later the molecularorbital picture has become the most common one, but there is still no consensus among chemists that it is always the best way to look at molecules. The reason is that this model of molecules and all others are based on simplifying assumptions and are thus approximate, partial descriptions. In reality, a molecule is a bunch of atomic nuclei in a cloud of electrons, with opposing electrostatic forces fighting a constant tug-of-war with one another, and all components constantly moving and reshuffling. Existing models of the molecule usually try to crystallize such a dynamic entity into a static one and may capture some of its salient properties but neglect others.
Quantum theory is unable to supply a unique definition of chemical bonds that accords with the intuition of chemists whose daily business is to make and break them. There are now many ways of describing molecules as atoms joined by bonds. According to quantum chemist Dominik Marx of Ruhr University Bochum in Germany, pretty much all such descriptions “are useful in some cases but fail in others.”
Computer simulations can now calculate the structures and properties of molecules from quantum first principles with great accuracy—as long as the number of electrons is relatively small. “Computational chemistry can be pushed to the level of utmost realism and complexity,” Marx says. As a result, computer calculations can increasingly be regarded as a kind of virtual experiment that predicts the course of a reaction. Once the reaction to be simulated involves more than a few dozen electrons, however, the calculations quickly begin to overwhelm even the most powerful supercomputer, so the challenge will be to see whether the simulations can scale up—whether, for example, complicated biomolecular processes in the cell or sophisticated materials can be modeled this way.
In the 1920s physicists Walter Heitler and Fritz London showed how to describe a chemical bond using the equations of then nascent quantum theory, and the great American chemist Linus Pauling proposed that bonds form when the electron orbitals of different atoms overlap in space. A competing theory by Robert Mulliken and Friedrich Hund suggested that bonds are the result of atomic orbitals merging into “molecular orbitals” that extend over more than one atom. Theoretical chemistry seemed about to become a branch of physics.
Nearly 100 years later the molecularorbital picture has become the most common one, but there is still no consensus among chemists that it is always the best way to look at molecules. The reason is that this model of molecules and all others are based on simplifying assumptions and are thus approximate, partial descriptions. In reality, a molecule is a bunch of atomic nuclei in a cloud of electrons, with opposing electrostatic forces fighting a constant tug-of-war with one another, and all components constantly moving and reshuffling. Existing models of the molecule usually try to crystallize such a dynamic entity into a static one and may capture some of its salient properties but neglect others.
Quantum theory is unable to supply a unique definition of chemical bonds that accords with the intuition of chemists whose daily business is to make and break them. There are now many ways of describing molecules as atoms joined by bonds. According to quantum chemist Dominik Marx of Ruhr University Bochum in Germany, pretty much all such descriptions “are useful in some cases but fail in others.”
Computer simulations can now calculate the structures and properties of molecules from quantum first principles with great accuracy—as long as the number of electrons is relatively small. “Computational chemistry can be pushed to the level of utmost realism and complexity,” Marx says. As a result, computer calculations can increasingly be regarded as a kind of virtual experiment that predicts the course of a reaction. Once the reaction to be simulated involves more than a few dozen electrons, however, the calculations quickly begin to overwhelm even the most powerful supercomputer, so the challenge will be to see whether the simulations can scale up—whether, for example, complicated biomolecular processes in the cell or sophisticated materials can be modeled this way.
SOURCE : SCIENTIFIC AMERICAN OCTOBER 2011
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