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No one can see an electron. Even an electron microscope, used for imaging the activities of these subatomic particles, does not offer a glimpse of an electron as one can look at an amoeba; instead, the microscope detects the patterns of electron deflection. In any case, a single-cell organism is gargantuan in comparison to an electron. Even when compared to a proton or a neutron, particles at the center of an atom, electrons are minuscule, being slightly more than 1/2000 the size of either. Yet the electron is the key to understanding the chemical process of bonding, and electron configurations clarify a number of aspects of the periodic table that may, at first, seem confusing.


The Electron's Place in the Atom

The atom is discussed in detail elsewhere in this book; here, the particulars of atomic structure will be presented in an abbreviated form, so that the discussion of electrons may proceed. At the center of an atom, the smallest particle of an element, there is a nucleus, which contains protons and neutrons. Protons are positive in charge, while neutrons exert no charge.

Moving around the nucleus are electrons, negatively charged particles whose mass is very small in comparison to the proton or neutron: 1/1836 and 1/1839 the mass of the proton and neutron respectively. The mass of an electron is 9.109389 · 1033 g. Compare the number 9 to the number 1,000,000,000,000,000,000,000,000,000,000,000,

and this gives some idea of the ratio between an electron's mass and a gram. The large number1 followed by 33 zeroesis many trillions of times longer than the age of the universe in seconds.


Electrons, though very small, are exceedingly powerful, and they are critical to both physical and chemical processes. The negative charge of an electron, designated by the symbol 1 or 1, is enough to counteract the positive charge (1+ or +1) of a proton, even though the proton has much greater mass. As a result, atoms (because they possess equal numbers of protons and electrons) have an electric charge of zero.

Sometimes an atom will release an electron, or a released electron will work its way into the structure of another atom. In either case, an atom that formerly had no electric charge acquires one, becoming an ion. In the first of the instances described, the atom that has lost an electron or electrons becomes a positively charged ion, or cation. In the second instance, an atom that gains an electron or electrons becomes a negatively charged ion, or anion.

One of the first forms of chemical bonding discovered was ionic bonding, in which electrons clearly play a role. In fact, electrons are critical to virtually all forms of chemical bonding, which relates to the interactions between electrons of different atoms.


If the size of the nucleus were compared to a grape, the edge of the atom itself would form a radius of about a mile. Because it is much smaller than the nucleus, an electron might be depicted in this scenario as a speck of dustbut a speck of dust with incredible energy, which crosses the mile radius of this enlarged atom at amazing speeds. In an item of matter that has been frozen to absolute zero, an electron moves relatively little. On the other hand, it can attain speeds comparable to that of the speed of light: 186,000 mi (299,339 km) per second.

The electron does not move around the nucleus as a planet orbits the Suna model of electron behavior that was once accepted, but which has since been overturned. On the other hand, as we shall see, the electron does not simply "go where it pleases": it acts in accordance with complex patterns described by the quantum theory of physics. Indeed, to some extent the behavior of the electron is so apparently erratic that the word "pattern" seems hardly to describe it. However, to understand the electron clearly, one has to set aside all ideas about how objects behave in the physical world.

Emerging Models of Electron Behavior

The idea that matter is composed of atoms originated in ancient Greece, but did not take hold until early in the nineteenth century, with the atomic theory of English chemist John Dalton (1766-1844). In the years that followed, numerous figuresamong them Russian chemist Dmitri Mendeleev (1834-1907), father of the periodic table of elementscontributed to the emerging understanding of the atomic model.

Yet Mendeleev, despite his awareness that the mass of an atom differentiated one element from another, had no concept of subatomic particles. No one did: until very late in the nineteenth century, the atom might as well have been a hard-shelled ball of matter, for all that scientists understood about its internal structure. The electron was the first subatomic particle discovered, in 1897nearly a century after the scientific beginning of atomic theory. Yet the first hints regarding its existence had begun to appear some 60 years before.


In 1838, British physicist and chemist Michael Faraday (1791-1867) was working with a set of electrodesmetal plates used to emit or collect electric chargewhich he had placed at either end of an evacuated glass tube. (In other words, most of the air and other matter had been removed from the tube.) He applied a charge of several thousand volts between the electrodes, and discovered that an electric current flowed between them.

This seemed to suggest the existence of particles carrying an electric charge, and four decades later, English physicist William Crookes (1832-1919) expanded on these findings with his experiments using an apparatus that came to be known as a Crookes tube. As with Faraday's device, the Crookes tube used an evacuated glass tube encased between two electrodesa cathode at the negatively charged end, and an anode at the positively charged end. A wire led outside the bulb to an electric source, and when electricity was applied to the electrodes, the cathodes emitted rays. Crookes concluded that the cathode rays were particles with a negative electric charge that came from the metal in the cathode plate.

English physicist J. J. Thomson (1856-1940) hypothesized that the negatively charged particles Crookes had observed were being emitted by atoms, and in 1897 he gave a name to these particles: electrons. The discovery of the electron raised questions concerning its place in the atom: obviously, there had to be a counterbalancing positive charge, and if so, from whence did it come?


Around the beginning of the twentieth century, the prevailing explanation of atomic structure was the "plum pudding model," which depicted electrons as floating like raisins in a "pudding" of positive charges. This was overturned by the discovery of the nucleus, and, subsequently, of the proton and neutron it contained.

In 1911, the great Danish physicist Niels Bohr (1885-1962) studied hydrogen atoms, and concluded that electrons move around the nucleus in much the same way that planets move around the Sun. This worked well when describing the behavior of hydrogen, which, in its simplest formthe isotope protiumhas only one electron and one proton, without any neutrons. The model did not work as well when applied to other elements, however, and within less than two decades Bohr's planetary model was overturned.

As it turns out, the paths of an electron's movement around the nucleus are nothing like that of a planet's orbitexcept inasmuch as both models describe a relatively small object moving around a relatively large one. The reality is much more complex, and to comprehend the secret of the electron's apparently random behavior is truly a mind-expanding experience. Yet Bohr is still considered among the greatest scientists of the twentieth century: it was he, after all, who first explained the quantum behavior of electrons, examined below.


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.


The wave mechanical model depicted the movement of electrons, not as smooth orbits, but as orbitalsregions 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 centermuch 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.


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.


Given the challenges involved in understanding electron behavior, it is amazing just how much scientists do know about electronsparticularly 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.


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 s1. 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, 2s (for now we will dispense with the superscript 1) and 2p. 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 2p x, 2p y, or 2p 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 sublevelss, 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.


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 s2. 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


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 2s is filled; from neon (atomic number 10), sublevel 2p and hence principal level 2 as a wholeis filled, and so on. The electron configuration for neon, thus, is written as 1s22 s22 p6.

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.


After argon, however, as one moves to the element occupying the nineteenth position on the periodic tablepotassiumthe rules change. Argon has an electron configuration of 1s 22s 22p 63s 23p 6, and by the pattern established with the first 18 elements, potassium should begin filling principal level 3d. Instead, it "skips" 3d and moves on to 4s. The element following argon, calcium, adds a second electron to the 4s 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 3d 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 4f orbitals are filled, while the actinide series reflects the filling of the seven 5f 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 factorsrelating to electrons, if not necessarily electron configurationthat explain why one element bonds as it does to another.


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 factthe increase of atomic size in a group, as a function of increasing atomic numberis 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.


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"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).

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



The negative ion that results when an atom gains one or more electrons.


An electrode at the positively charged end of a supply of electric current.


The smallest particle of an element.


The number of protons in the nucleus of an atom. Since this number is different for each element, elements are listed on the periodic table of elements in order of atomic number.


An electrode at the negatively charged end of a supply of electric current.


The positive ion that results when an atom loses one or more electrons.


A structure, often a metal plate or grid, that conducts electricity, and which is used to emit or collect electriccharge.


A negatively charged particle in an atom.


A term used to describe the pattern formed by orbitals.


A term describing thecharacteristics of an atom that has acquired excess energy.


A term describing the state of an atom at its ordinary energy level.


An atom or atoms that has lost or gained one or more electrons, and thus has a net electric charge.


Atoms that have an equal number of protons, and hence are of the same element, but differ in their number of neutrons.


A subatomic particle that has no electric charge. Neutrons are found in the nucleus of an atom, alongside protons.


The center of an atom, a region where protons and neutrons are located, and around which electrons spin.


A pattern of probabilities regarding the regions that an electron can occupy within an atom in a particular energy state. The orbital, complex and imprecise as it may seem, is a much more accurate depiction of electron behavior than the model once used, which depicted electrons moving in precisely defined orbits around the nucleus, rather as planets move around the Sun.


A chart that shows the elements arranged in order of atomic number. Vertical columns within the periodic table indicate groups or "families" of elements with similar chemical characteristics.


A particle of electromagnetic radiation.


A value indicating the distance that an electron may move away from the nucleus of anatom. 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.


A positively charged particle in an atom.


A term describing any property that has only certain discrete values, as opposed to values distributed along a continuum. The quantization of an atom means that it does not have a continuous range of energy levels; rather, it can exist only at certain levels of energy from the ground state through various excitedstates.


In a general sense, radiation can refer to anything that travels in astream, whether that stream be composed of subatomic particles or electromagnetic waves.

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