Lewis Structure

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Lewis Structure


Lewis structures (also called electron-dot structures) are formed when Lewis symbols (also called electron-dot symbols) are combined. Lewis symbols are a simple way of visualizing the valence electrons in an atom. In a Lewis symbol, the symbol for the element is used to represent the atom and its core electrons. Dots placed around the atom are used to indicate the valence electrons. When combined to form Lewis structures, Lewis symbols make it possible to predict the shape of many molecules and ions. This information is particularly useful as many physical

and chemical properties of molecules and ions are dependent on their shape.

Subsequent to the discovery of the nuclear atom and electrons, many theories were proposed to explain the nature of chemical bonding in molecules using electrons. In 1916, Gilbert Newton Lewis, an American chemist, suggested that molecules were formed when atoms shared pairs of outer electrons. He also made the assumption that unshared electrons were found in pairs and proposed that the arrangement of eight electrons around an atom was a particularly stable configuration. In 1919, another American chemist named Irving Langmuir, noted that Lewiss proposed theory accounted particularly well for the chemical facts known at that time. He suggested the name covalent bond for a pair of shared electrons. Lewiss theory, although very useful, did not explain why or how electrons were shared, however, and covalent bonds were not well understood until quantum theory was developed.

In a neutral atom, the number of protons in the nucleus equals the number of electrons around the nucleus. The electrons are assigned energy levels based on quantum theory. The closer the electrons are to the nucleus, the more tightly they are bound. These tightly bound electrons are difficult to remove from the atom and are called the inner-core electrons. The electrons in the outermost layer (those that are furthest from the nucleus) are most loosely bound.

Think of electrons in an atom like students in an auditorium. The students in the rows near the front are closest to the lecturer. It is difficult for these students to misbehave or fall asleep as the lecturer can keep a close eye on them. As students sit farther and farther away from the lecturer, it is more difficult for the lecturer to notice what they are doing and it is even possible for a student at the back of the auditorium to leave without the lecturer noticing. Similarly, as electrons are placed in energy levels farther and farther from the nucleus, it is easier for them to leave the atom or to react. The electrons in the outermost energy level are called valence electrons. According to quantum mechanics, the outermost level can hold at most eight electrons. In the periodic table, these electrons are filled (usually in order) in the first and second columns and then the last six columns. The first and second columns are called the s block and the last six columns are called the p block respectively, denoting the shape of the space most likely to contain these electrons.

In the Lewis symbol for an element in the s or p block, the number of valence electrons is indicated by dots placed around the symbol for the element. No electron is paired until four single electrons have been placed. Then, electrons are paired until eight electrons have been placed. The dots are usually placed on top, to the right, below, and to the left of the symbol. The order in which they are placed is unimportant. According to Lewis, the number of unpaired electrons indicated the number of bonds the atom usually formed in a compound. Using the second row of the periodic table as an example, lithium (having one unpaired electron) would form one bond, beryllium (having two unpaired electrons) would form two bonds, boron (having three unpaired electrons) would form three bonds, and carbon (having four unpaired electrons) would form four bonds.

After carbon, nitrogen has only three unpaired electrons (as the fifth electron forms a pair with one of the other electrons) and hence forms only three bonds. Similarly, oxygen forms only two bonds and fluorine forms only one bond. Neon (and all the elements in that last column) have eight electrons, all paired, and hence it was predicted that these elements would not form any bonds (a fact that was observed until very recently when some of the elements in the last column were forced to react under extreme conditions). This lack of reactivity on the part of the elements in the last column (called the inert or noble gases), led to the observation that it was highly desired to have eight electrons around the atom. The noble gases are said to have a stable octet configuration.

Lewis structures are formed when two or more Lewis symbols are combined. One unshared electron from each atom combines to form a pair of electrons shared between two atoms. This shared pair is called a covalent bond. Electrons in a covalent bond are also known as a bonding pair. If an atom has electrons that are in pairs but are not shared with another atom, these electron pairs are called lone pairs, unshared pairs, or nonbonding electrons. For example, hydrogen has one electron and chlorine has seven electrons in the form of three pairs and one lone electron. The lone electron of hydrogen can pair with the lone electron of chlorine to form a bonding pair of electrons


Bonding pair The pair of electrons shared between two atoms to form a bond is called a bonding pair.

Covalent bond A chemical bond formed when two atoms share a pair of electrons with each other.

Double bond A covalent bond consisting of two pairs of shared electrons that hold the two atoms together.

Free radical A species in which there is an unpaired electron. Free radicals are extremely reactive.

Lewis symbol Lewis symbols are combined to form the Lewis structure. In a Lewis symbol, the symbol of the element represents the nucleus and the core electrons. Dots placed around the symbol of the element represent the valence electrons.

Lone pair A pair of electrons located around an atom but not shared with another atom is called a lone pair, an unshared pair or a non-bonding pair.

Neutral Having neither a positive charge nor a negative charge.

Octet rule The octet rule states that all atoms, with the exception of hydrogen, strive to acquire eight electrons in their valence shell.

Stable octet configuration Any atom that has eight electrons in its valence shell is said to have a stable octet configuration.

Triple bond A triple bond is formed when three pairs of electrons are shared between two atoms.

Valence electrons The electrons in the outermost shell of an atom that determine an elements chemical properties.

and thus a covalent bond between hydrogen and chlorine. The chlorine, as a result, has one bonding pair and three lone pairs (a total of eight electrons), while hydrogen has one pair of bonding electrons. If Lewis structures are drawn for many other molecules, it can be seen that most atoms in stable molecules or ions, with the exception of hydrogen, acquire eight electrons in their outermost valence level. As mentioned previously, this is the stable octet configuration possessed by the noble gases. This observation led to the development of the octet rulethat each atom, with the exception of hydrogen, strives to acquire eight electrons in its valence shell. There are, however, a number of fairly common and important exceptions to this rule.

The abovementioned method is easy to follow when only two atoms are combined to make a molecule. When more than two atoms are combined, it is easier to follow a simple procedure when constructing the Lewis diagram. The first step is to write the atoms in the molecule in the order in which they are joined to one another. In general, hydrogen forms only one bond, and oxygen only bonds to another oxygen in peroxides.

Although there are exceptions, two generally accepted methods of writing the formula for a compound are used: The first is used when there is one atom with many other atoms attached to it. In this case, the central atom is usually written first, followed by the atoms that are attached to it. The second is used when there are multiple structures possible or when there is more than one central atom. In this case, the formula is written in the order in which the atoms are joined. Once the skeleton structure of the molecule is known, it is necessary to count the total number of valence electrons. This can be done by adding the number of valence electrons present in each element in the compound.

If the Lewis structure of an ion (a charged species) is being determined, it is also necessary to account for the charge of the ion. If the ion has a negative charge, a number of electrons equaling the numerical charge of the ion are added to the total number of electrons. If the ion has a positive charge, a number of electrons equaling the numerical charge of the ion are subtracted from the total number of electrons. Once the total number of electrons is known, the electrons are placed around the atoms in the Lewis structure. First, two electrons are placed between any two atoms that are joined to create a covalent bond or a bonding pair. Remaining electrons are distributed so that each atom acquires an octet (eight) of electrons.

In many cases, the above procedure provides a Lewis structure that can be used to determine much information. However, occasionally, there are insufficient electrons to provide each atom with an octet of electrons. If this is the case, it may be necessary to consider multiple bonding, in which two atoms share more than one pair of electrons. If two pairs of electrons are shared between atoms, the bond between the atoms is called a double bond. If three pairs are shared, the bond is called a triple bond.

If a molecule has double or triple bonds, alternate structures called resonance structures are possible; these involve different but viable arrangements of electrons between the atoms in the molecule or ion. It is important to note here that only the arrangement of the electrons differs. The arrangement of the atoms must stay the same. When there are a number of resonance structures, the true structure is none of the resonance structures, but is said to be a hybrid of the resonance structures.

Alternatively, there are many molecules that end up with an excess number of electrons. These are positioned around the central atom, which is usually sulfur or an element higher in atomic number. Finally, it is also possible, especially in the case where there is a nitrogen atom in the molecule, that there will be an odd number of electrons. In such cases, one electron remains unpaired. Such species are called radicals or free radicals. They are very reactive and have important roles in atmospheric chemistry, the chemistry of ageing, and in cancer.

Once the Lewis structure has been determined, it is possible to know the shape of the molecule or ion. The most important piece of information needed to determine the shape is the total number of groups around the central atom, where a group could be another atom or a lone pair. A central atom connected to one or two other atoms is linear. When the central atom is connected to three atoms, the shape is trigonal planar. When the central atom is connected to four atoms, the shape is tetrahedral. When the central atom is connected to five atoms the shape is trigonal bipyramidal (two triangular-based pyramids joined at the base). When the central atom is connected to six atoms, the shape is octahedral. Other shapes are possible when atoms are replaced with lone pairs.



Brady, Russell, and Holum. Chemistry, Matter and Its Changes. 3rd ed. New York: John Wiley and Sons Inc. 2000.

Zumdahl. Chemistry. 4th ed. New York: Houghton Mifflin 1997.


St.Olaf College, Chemistry Department. Construct a Lewis Structure <http://www.stolaf.edu/depts/chemistry/courses/toolkits/123/js/lewis/> (accessed December 2, 2006).

University of Waterloo, Faculty of Science. Lewis Dot Structures <http://www.science.uwaterloo.ca/cchieh/cact/c120/dotstruc.html> (accessed December 2, 2006).

Rashmi Venkateswaran

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Lewis Structure

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