Electrons - Real-life applications



Quantum Theory and the Atom

Much of what scientists understand today about the atom in general, and the electron in particular, comes from the quantum theory introduced by German physicist Max Planck (1858-1947). Planck showed that, at the atomic level, energy is emitted in tiny packets, or "quanta." Applying this idea to the electron, Bohr developed an idea of the levels at which an electron moves around the nucleus. Though his conclusions led him to the erroneous planetary model, Bohr's explanation of energy levels still prevails.

As has been suggested, the interaction between electrons and protons is electromagnetic, and electromagnetic energy is emitted in the form of radiation, or a stream of waves and particles. The Sun, for instance, emits electromagnetic radiation along a broad spectrum that includes radio waves, infrared light, visible light, ultraviolet light, x rays, and gamma rays. These are listed in ascending order of their energy levels, and the energy emitted can be analyzed in terms of wavelength and frequency: the shorter the wavelength, the greater the frequency and the greater the energy level.

When an atom is at its ordinary energy level, it is said to be in a ground state, but when it acquires excess energy, it is referred to as being in an excited state. It may release some of that energy in the form of a photon, a particle of electromagnetic radiation. The amount of energy involved can be analyzed in terms of the wavelengths of light the atom emits in the form of photons, and such analysis reveals some surprising things about the energy levels of atoms.

If one studies the photons emitted by an atom as it moves between a ground state and an excited state, one discovers that it emits only certain kinds of photons. From this, Bohr concluded that the energy levels of an atom do not exist on a continuum; rather, there are only certain energy levels possible for an atom of a given element. The energy levels are therefore said to be quantized.

The Wave Mechanical Model

In everyday terms, quantization can be compared to the way that a person moves up a set of stairs: by discrete steps. If one step directly follows another, there is no step in between, nor is there any gradual way of moving from step to step, as one would move up a ramp. The movement of electrons from one energy level to another is not a steady progression, like the movement of a person up a ramp; rather, it is a series of quantum steps, like those a person makes when climbing a set of stairs.

The idea of quantization was ultimately applied to describing the paths that an electron makes around the nucleus, but this required some clarification along the way. It had been believed that an electron could move through any point between the nucleus and the edge of the atom (again, like a ramp), but it later became clear that the electrons could only move along specific energy levels. As we have seen, Bohr believed that these corresponded to the orbits of planets around the Sun; but this explanation would be discarded in light of new ideas that emerged in the 1920s.

During the early part of that decade, French physicist Louis de Broglie (1892-1987) and Austrian physicist Erwin Schrödinger (1887-1961) introduced what came to be known as the wave mechanical model, also known as the particle-wave hypothesis. Because light appeared to have the properties of both particles and waves, they reasoned, electrons (possessing electromagnetic energy as they did) might behave in the same fashion. In other words, electrons were not just particles: in some sense, they were waves as well.

UNDERSTANDING ORBITALS.

The wave mechanical model depicted the movement of electrons, not as smooth orbits, but as orbitals—regions in which there is the highest probability that an electron will be found. An orbital is nothing like the shape of a solar system, but, perhaps ironically, it can be compared to the photographs astronomers have taken of galaxies. In most of these photographs, one sees an area of intense light emitted by the stars in the center. Further from this high-energy region, the distribution of stars (and hence of light) becomes increasingly less dense as one moves from the center of the galaxy to the edges.

Replace the center of the galaxy with the nucleus of the atom, and the stars with electrons, and this is an approximation of an orbital. Just as a galaxy looks like a cloud of stars, scientists use the term electron cloud to describe the pattern formed by orbitals. The positions of electrons cannot be predicted; rather, it is only possible to assign probabilities as to where they will be. Naturally, they are most drawn to the positive charges in the nucleus, and hence an orbital depicts a high-density region of probabilities at the center—much like the very bright center of a galaxy.

The further away from the nucleus, the less the probability that an electron will be in that position. Hence in models of an orbital, the dots are concentrated at the center, and become less dense the further away from the nucleus they are. As befits the comparison to a cloud, the edge of an orbital is fuzzy. Contrary to the earlier belief that an atom was a clearly defined little pellet of matter, there is no certainty regarding the exact edge of a given atom; rather, scientists define the sphere of the atom as the region encompassing 90% of the total electron probability.

THE MYSTERIOUS ELECTRON.

As complex as this description of electron behavior may seem, one can rest assured that the reality is infinitely more complex than this simplified explanation suggests. Among the great mysteries of the universe is the question of why an electron moves as it does, or even exactly how it does so. Nor does a probability model give us any way of knowing when an electron will occupy a particular position.

In fact, as German physicist Werner Heisenberg (1901-1976) showed with what came to be known as the Heisenberg Uncertainty Principle, it is impossible to know both the speed of an electron and its precise position at the same time. This, of course, goes against every law of physics that prevailed until about 1920, and in fact quantum theory offers an entirely different model of reality than the one accepted during the seventeenth, eighteenth, and nineteenth centuries.

Orbitals

Given the challenges involved in understanding electron behavior, it is amazing just how much scientists do know about electrons—particularly where energy levels are concerned. This, in turn, makes possible an understanding of the periodic table that would astound Mendeleev.

Every element has a specific configuration of energy levels that becomes increasingly complex as one moves along the periodic table. In the present context, these configurations will be explained as simply as possible, but the reader is encouraged to consult a reliable chemistry textbook for a more detailed explanation.

PRINCIPAL ENERGY LEVELS AND SUBLEVELS.

The principal energy level of an atom indicates a distance that an electron may move away from the nucleus. This is designated by a whole-number integer, beginning with 1 and moving upward: the higher the number, the further the electron is from the nucleus, and hence the greater the energy in the atom. Each principal energy level is divided into sublevels corresponding to the number n of the principal energy level: thus principal energy level 1 has one sublevel, principal energy level 2 has two, and so on.

The simplest imaginable atom, a hydrogen atom in a ground state, has an orbital designated as 1 s 1 . The s indicates that an electron at energy level 1 can be located in a region described by a sphere. As for the significance of the superscript 1, this will be explained shortly.

Suppose the hydrogen atom is excited enough to be elevated to principal energy level 2. Now there are two sublevels, 2 s (for now we will dispense with the superscript 1) and 2 p. A p orbital is rather like the shape of a figure eight, with its center of gravity located on the nucleus, and thus unlike the s sublevel, p orbitals can have a specific directional orientation. Depending on whether it is oriented along an x− , y− , or z− axis, orbitals in sublevel p are designated as 2 p x , 2 p y , or 2 p z .

If the hydrogen atom is further excited, and therefore raised to principal energy level 3, it now has three possible sublevels, designated as s , p , and d. Some of the d orbitals can be imagined as two figure eights at right angles to one another, once again with their centers of gravity along the nucleus of the atom. Because of their more complex shape, there are five possible spatial orientations for orbitals at the d sublevel.

Even more complex is the model of an atom at principal level 4, with four sublevels— s , p , d , and f , which has a total of seven spatial orientations. Obviously, things get very, very complex at increased energy levels. The greater the energy level, the further the electron can move from the nucleus, and hence the greater the possible number of orbitals and corresponding shapes.

ELECTRON SPIN AND THE PAULI EXCLUSION PRINCIPLE.

Every electron spins in one of two directions, and these are indicated by the symbols ↑ and ↓. According to the Pauli exclusion principle, named after the Austrian-Swiss physicist Wolfgang Pauli (1900-1958), no more than two electrons can occupy the same orbital, and those two electrons must have spins opposite one another.

This explains the use of the superscript 1, which indicates the number of electrons in a given orbital. This number is never greater than two: hence, the electron configuration of helium is written as 1 s 2 . It is understood that these two electrons must be spinning in opposite directions, but sometimes this is indicated by an orbital diagram showing both an upward-and downward-pointing arrow in an orbital that has been filled, or only an upward-pointing arrow in an orbital possessing just one electron.

Electron Configuration and the Periodic Table

THE FIRST 18 ELEMENTS.

As one moves up the periodic table from atomic number 1 (hydrogen) to 18 (argon), a regular pattern emerges. The orbitals are filled in a neat progression: from helium (atomic number 2) onward, all of principal level 1 is filled; beginning with beryllium (atomic number 4), sublevel 2 s is filled; from neon (atomic number 10), sublevel 2 p —and hence principal level 2 as a whole—is filled, and so on. The electron configuration for neon, thus, is written as 1 s 2 2 s 2 2 p 6 .

Note that if one adds together all the superscript numbers, one obtains the atomic number of neon. This is appropriate, of course, since atomic number is defined by the number of protons, and an atom in a non-ionized state has an equal number of protons and electrons. In noticing the electron configurations of an element, pay close attention to the last or highest principal energy level represented. These are the valence electrons, the ones involved in chemical bonding. By contrast, the core electrons, or the ones that are at lower energy levels, play no role in the bonding of atoms.

SHIFTS IN ELECTRON CONFIGURATION PATTERNS.

After argon, however, as one moves to the element occupying the nineteenth position on the periodic table—potassium—the rules change. Argon has an electron configuration of 1 s 2 2 s 2 2 p 6 3 s 2 3 p 6 , and by the pattern established with the first 18 elements, potassium should begin filling principal level 3d. Instead, it "skips" 3 d and moves on to 4 s. The element following argon, calcium, adds a second electron to the 4 s level.

After calcium, the pattern again changes. Scandium (atomic number 21) is the first of the transition metals, a group of elements on the periodic table in which the 3 d orbitals are filled. This explains why the transition metals are indicated by a shading separating them from the rest of the elements on the periodic table.

But what about the two rows at the very bottom of the chart, representing groups of elements that are completely set apart from the periodic table? These are the lanthanide and actinide series, which are the only elements that involve f sublevels. In the lanthanide series, the seven 4 f orbitals are filled, while the actinide series reflects the filling of the seven 5 f orbitals.

As noted earlier, the patterns involved in the f sublevel are ultra-complex. Thus it is not surprising than the members of the lanthanide series, with their intricately configured valence electrons, were very difficult to extract from one another, and from other elements: hence their old designation as the "rare earth metals." However, there are a number of other factors—relating to electrons, if not necessarily electron configuration—that explain why one element bonds as it does to another.

CHANGES IN ATOMIC SIZE.

With the tools provided by the basic discussion of electrons presented in this essay, the reader is encouraged to consult the essays on Chemical Bonding, as well as The Periodic Table of the Elements, both of which explore the consequences of electron arrangement in chemistry. Not only are electrons the key to chemical bonding, understanding their configurations is critical to an understanding of the periodic table.

One of the curious things about the periodic table, for instance, is the fact that the sizes of atoms decrease as one moves from left to right across a row or period, even though the sizes increase as one moves from top to bottom along a column or group. The latter fact—the increase of atomic size in a group, as a function of increasing atomic number—is easy enough to explain: the higher the atomic number, the higher the principal energy level, and the greater the distance from the nucleus to the furthest probability range.

On the other hand, the decrease in size across a period (row) is a bit more challenging to comprehend. However, all the elements in a period have their outermost electrons at a particular principal energy level corresponding to the number of the period. For instance, the elements on period 5 all have principal energy levels 1 through 5. Yet as one moves along a period from left to right, there is a corresponding increase in the number of protons within the nucleus. This means a stronger positive charge pulling the electrons inward; therefore, the "electron cloud" is drawn ever closer toward the increasingly powerful charge at the center of the atom.

WHERE TO LEARN MORE

"Chemical Bond" (Web site). <http://www.science.uwaterloo.ca/~cchieh/cact/c120/chembond.html> (May 18, 2001).

"The Discovery of the Electron" (Web site). <http://www.aip.org/history/electron/> (May 18, 2001).

Ebbing, Darrell D.; R. A. D. Wentworth; and James P. Birk. Introductory Chemistry. Boston: Houghton Mifflin, 1995.

Gallant, Roy A. The Ever-Changing Atom. New York: Benchmark Books, 1999.

Goldstein, Natalie. The Nature of the Atom. New York: Rosen Publishing Group, 2001.

"Life, the Universe, and the Electron" (Web site). <http://www.iop.org/Physics/Electron/Exhibition/> (May 18, 2001).

"A Look Inside the Atom" (Web site). <http://www.aip.org/history/electron/jjhome.htm> (May 18, 2001).

"Valence Shell Electron Pair Repulsion (VSEPR)." <http://www.shef.ac.uk/~chem/vsepr/chime/vsepr.html> (May 18, 2001).

"What Are Electron Microscopes?" (Web site). <http://www.unl.edu/CMRAcfem/em.htm> (May 18, 2001).

Zumdahl, Steven S. Introductory Chemistry: A Foundation, 4th ed. Boston: Houghton Mifflin, 2000.



User Contributions:

Comment about this article, ask questions, or add new information about this topic: