Structure Lewis

Lewis early depiction of the valence shell of the period 2 elements. 

Lewis structures for (a) methane, (b) ethylene, and (c) ethyne, where the electrons are positioned at the corners of a tetrahedron for each carbon atom. 

While the precise definition of a Lewis structure has changed somewhat over the years, the essence of the concept is virtually the same. Lewis structures remain one of the single most powerful pedagogical tools in the understanding of chemical bonding. Modern Lewis structures for simple molecules can be constructed from the following set of rules. 

If the FC on the central atom is zero or equal to the charge of the polyatomic ion, then the structure is correct as it is. 

For Series 2 elements, continue this process until the central atom has a full octet. For all elements higher than Series 2, continue this process until the central atom has an FC of zero. 

Lewis developed a notation that allows us to use the valence electrons of atoms in a molecule to predict the bonding in that molecule. In this method, the electrons in the valence shell of each atom are drawn as dots for all atoms in the molecule (see examples below). Bonds are formed by sharing of one or more pairs of electrons between the atoms, such that each atom achieves an octet of electrons. In an alternative to the electron dot symbolism, we can draw a line to represent a bond. A single bond is the sharing of two electrons, while double and triple bonds involve the sharing of four and six electrons, respectively. Despite its simplicity, this notation can be used to accurately predict the number of lone pairs that an atom will have and whether that atom will use single, double, or triple bonds when incorporated into specific molecules. 

The problem with Lewis dot structures is that they provide no insight into molecular shapes, orbitals, or distributions of electrons within molecules. Instead, they are only useful for predicting the number of bonds an atom forms whether the atom has lone pairs and whether single, double, or triple bonds are used. Once an atom is found to have an octet using a Lewis analysis, no further insight into the structure or reactivity can be obtained from the Lewis structure. We have to turn to more sophisticated molecular structure and bonding concepts to understand structure and reactivity. 

Lewis structures are electron dot representations for molecules. There are three general rules for drawing Lewis structures. 

Give every second-row element an octet of electrons, if possible. 

To draw a Lewis structure for a diatomic molecule like HF, recall that hydrogen has one valence electron and fluorine has seven. H and F each donate one electron to form a two-electron bond. The resulting molecule gives both H and F a filled valence shell. 

In a Lewis structure, a solid line indicates a two-electron covalent bond. 

Summary The usual number of bonds of common neutral atoms 

Lewis structures show only valence electrons. 

Lewis (above) conceived the octet rule while lecturing to a class of general chemistry students in 1902. He was also one of the two authors of a now classic work on thermodynamic, Lewis and Randall, Thermodynamics and the Free Energy of Chemical Substances (1923). (right) This is his original sketch. From G. N. Lewis, Valence, Dover Publications, Inc., New York, 1966. 

Description of the valence electron arrangement in the molecule using Lewis structures (will he discussed in the next section). 

Prediction of the geometry of the molecule, using the valence shell electron-pair repulsion (VSEPR) model (will be discussed in Section 13.13). 

Description of the types of atomic orbitals used by the atoms to share electrons or hold lone pairs (will be discussed in Chapter 14). 

These representations are all Lewis structures, formulas in which atomic symbols represent nuclei and inner-shell electrons, dot-pairs or dashes between two atomic symbols represent electron pairs in covalent bonds, and dots adjacent to only one atomic symbol represent unshared electrons. It is 

The Lewis structures (and therefore the structural formulas) for many molecules can be drawn if one knows the composition of the molecule and which atoms are bonded to each other. The following sample problem illustrates the basic steps for writing Lewis structures. The molecule described in this problem contains bonds with single, shared electron pairs. A singie covaient bond, or a single bond, is a covalent bond in which one pair of electrons is shared between two atoms. 

Sample Problem C Draw the Lewis structure of iodomethane,CH3l.  

0 SOLVE 1. Determine the type and number of atoms in the molecule. The formula shows one carbon atom, one iodine atom, and three hydrogen atoms. 2. Write the electron-dot notation for each type of atom in the molecule. Carbon is from Group 14 and has four valence electrons. Iodine is from Group 17 and has seven valence electrons. Hydrogen has one valence electron. 

Determine the total number of valence electrons available in the atoms to be combined. 

In the early part of the last century, a chemist at the University of California at Berkeley named Gilbert N. Lewis devised a system for diagramming atoms and molecules. Though simple, the system is still used today because predictions made from these diagrams often match those based on experiment. 

Lewis proposed the following representations for the first ten elements with their valence electrons. 

Only He and Ne are found in nature as shown above. All the other elements are found either as a charged species (ion) or as part of a molecule that can be represented as a legitimate Lewis structure. 

CHECKLIST a Legitimate Lewis Structure is a dot or line bond representation in which... 

The correct TOTAL number of valence electrons is shown. 

The Lewis stmctnre of a molecule shows how the valence electrons are arranged among the atoms in the molecule. These representations are named after G. N. Lewis (Eig. 8.14). The rules for writing Lewis structures are based on observations of thousands of molecules. Erom experiment, chemists have learned that the most important requirement for the formation of a stable compound is that the atoms achieve noble gas electron configurations. 

Unless otherwise noted, all arton this page is Cengage Learning2014. 

We need to learn the number of bonds that are commonly formed by the atoms that are important in organic chemistry. In the vast majority of cases, hydrogen forms one covalent bond so that it has two electrons in its outer shell. In covalent compounds the atoms of the second period (or row) of the periodic table need eight electrons in their outer shell. (This is called the octet rule.) We have already seen that to accomplish this carbon forms four bonds. In a similar fashion we can determine the number of bonds that some other second-period atoms prefer to form. The number of bonds that commonly occur in neutral molecules for the elements of most interest in organic chemistry are listed in Table 1.1. 

Show Lewis structures for the simplest neutral compounds formed from these elements a) Carbon and chlorine b) Hydrogen and bromine 

Although the number of bonds shown in Table 1.1 yields neutral structures and is the most common number for each of these elements, other bonding arrangements are also encountered. For example, in some cases, oxygen forms a third bond by using a 

Show a Lewis structure for the simplest neutral compound formed from hydrogen and sulfur. 

It is useful to be able to look at a Lewis structure and judge the stability of the compound represented by that structure. Can the compound be isolated and put in a bottle Is it more stable than another, similar compound The most important factor to examine in determining whether a molecule is stable is whether the octet rule is satisfied. Compounds with fewer than eight electrons around a second-period atom are known but are seldom stable. (However, such compounds are encountered as unstable, highly reactive intermediates in some chemical reactions.) Furthermore, compounds with more 

Covalent bonding, in which electrons are shared rather than transferred, is the most common type of bonding in organic compounds. Hydrogen, for example, needs a second electron to achieve the noble-gas configuration of helium. If two hydrogen atoms come together and form a bond, they share their two electrons, and each atom has two electrons in its valence shell. 

We will study covalent bonding in more detail in Chapter 2. 

One way to symbolize the bonding in a covalent molecule is to use Lewis structures. In a Lewis structure, each valence electron is symbolized by a dot A bonding pair of electrons is symbolized by a pair of dots or by a dash (— ). We try to arrange all the atoms so they have their appropriate noble-gas configurations two electrons for hydrogen, and octets for the second-row elements. 

Carbon contributes four valence electrons, and each hydrogen contributes one, to give a total of eight electrons. All eight electrons surround carbon to give it an octet, and each hydrogen atom shares two of the electrons. 

The Lewis structure for ethane (C2Hg) is more complex. 

Remember that the electrons in the highest principal energy level of an atom are called the valence electrons. 

Module 12 Drawing Lewis Electron Dot Structures covers concepts in this section, 

Bonding involves just the valence electrons of atoms. Valence electrons are transferred when a metal and a nonmetal react to form an ionic compound. Valence electrons are shared between nonmetals in covalent bonds. 

We have already seen this rule operate in the reaction of metals and nonmetals to form binary ionic compounds. An example is the formation of KBr, where the ion has the [Ar] electron configuration and the Br ion has the [Kr] electron configuration. In writing Lewis structures, we include only the valence electrons. Using dots to represent valence electrons, we write the Lewis structure for KBr as follows  

1 In which of the following molecules are the bonds most polar  

2 Using data from Table 8.5, calculate the magnitude of the partial charges in HI. 

Arrange molecules A through E in order of increasing percent ionic character. 

Although the octet rule and Lewis structures alone do not present a complete picture of covalent bonding, they do help us account for some of the properties of molecules. In addition, Lewis stmctures provide a starting point for the bonding theories that we will examine in Chapter 9. It is cmcial, therefore, that you leam a system for drawing correct Lewis stmctures for molecules and polyatomic ions. The basic steps are as follows  

Student Annotation H cannot be a central atom because it only forms one covalent bond 

No dots are shown on the K ion because it has lost its only valence electron (the 4s electron). The Br ion is shown with eight electrons because it has a filled valence shell. 

Next we will consider Lewis structures for molecules with covalent bonds, involving nonmetals in the first and second periods. The principle of achieving a noble gas electron configuration applies to these elements as follows  

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Structural formulas of this type in which electrons are represented as dots are called Lewis structures... 

Representing a two electron covalent bond by a dash (—) the Lewis structures for hydrogen fluoride fluorine methane and carbon tetrafluoride become... 

Lewis s concept of shared electron parr bonds allows for four electron double bonds and SIX electron triple bonds Carbon dioxide (CO2) has two carbon-oxygen double bonds and the octet rule is satisfied for both carbon and oxygen Similarly the most stable Lewis structure for hydrogen cyanide (HCN) has a carbon-nitrogen triple bond... 

Multiple bonds are very common m organic chemistry Ethylene (C2H4) contains a carbon-carbon double bond m its most stable Lewis structure and each carbon has a completed octet The most stable Lewis structure for acetylene (C2H2) contains a carbon-carbon triple bond Here again the octet rule is satisfied... 

Write the most stable Lewis structure for each of the following... 

Lewis structures frequently contain atoms that bear a positive or negative charge If the molecule as a whole is neutral the sum of its positive charges must equal the sum of Its negative charges An example is nitric acid HNO3... 

Like nitric acid each of the following inorganic compounds will be frequently encountered in this text Calculate the formal charge on each of the atoms in the Lewis structures given... 

The formal charges are shown in the Lewis structure of thionyl chloride... 

Formal charges are based on Lewis structures m which electrons are considered to be shared equally between covalently bonded atoms Actually polarization of N—H bonds m ammonium ion and of B—H bonds m borohydride leads to some transfer of positive and negative charge respectively to the hydrogens... 

The structure shown is the best (most stable) Lewis structure for methyl nitrite All atoms except hydro gen have eight electrons (shared + unshared) in their valence shell... 

None of the atoms in the Lewis structure shown in step 6 possesses a formal charge An alternative Lewis structure for methyl nitrite... 

Table 1 4 summarizes the procedure we have developed for writing Lewis structures Notice that the process depends on knowing not only the molecular formula but also the order m which the atoms are attached to one another This order of attachment is called the constitution, or connectivity, of the molecule and is determined by experiment Only rarely is it possible to deduce the constitution of a molecule from its molecular formula... 

When writing a Lewis structure we restrict a molecule s electrons to certain well defined locations either linking two atoms by a covalent bond or as unshared electrons on a sm gle atom Sometimes more than one Lewis structure can be written for a molecule espe cially those that contain multiple bonds An example often cited m introductory chem istry courses is ozone (O3) Ozone occurs naturally m large quantities m the upper atmosphere where it screens the surface of the earth from much of the sun s ultraviolet rays Were it not for this ozone layer most forms of surface life on earth would be dam aged or even destroyed by the rays of the sun The following Lewis structure for ozone satisfies fhe ocfef rule all fhree oxygens have eighf elecfrons m fheir valence shell... 

Three equally stable Lewis structures are possible for nitrate ion The negative charge in nitrate is shared equally by all three oxygens... 

It IS good chemical practice to represent molecules by their most stable Lewis structure The ability to write alternative resonance forms and to compare their relative stabilities however can provide insight into both molecular structure and chemical behavior This will become particularly apparent m the last two thirds of this text where the resonance concept will be used regularly... 

Lewis structures in which second row elements own or share more than eight valence electrons are especially unstable and make no contribution to the true structure (The octet rule may be ex ceeded for elements beyond the second row)... 

The two Lewis structures D and E of methyl nitrite satisfy the octet rule... 

The most stable Lewis structure for cyanate ion is F because the negative charge is on its oxygen... 

Each contributing Lewis structure must have the same number of electrons and the same net charge although the formal charges of individual atoms may vary among the various Lewis struc tures... 

Each contributing Lewis structure must have the same number of unpaired electrons... 

Structural formula J is a Lewis structure of nitro methane K is not even though it has the same atomic positions and the same number of electrons... 

Electron delocalization stabilizes a molecule A molecule in which electrons are delocalized is more stable than implied by any of the individual Lewis structures that may be written for it The de gree of stabilization is greatest when the contrib uting Lewis structures are of equal stability... 

The opening paragraph of this chapter emphasized that the connection between structure and properties is what chemistry is all about We have just seen one such con nection From the Lewis structure of a molecule we can use electronegativity to tell us about the polarity of bonds and combine that with VSEPR to predict whether the mol ecule has a dipole moment In the next several sections we 11 see a connection between structure and chemical reactivity as we review acids and bases... 

Electron Delocalization in the Conjugate Base With a of —1 4 nitnc acid is almost completely ionized m water If we look at the Lewis structure of nitric acid m light of what we have said about inductive effects we can see why The N atom m nitric acid IS not only electronegative m its own right but bears a formal charge of +1 which enhances its ability to attract electrons away from the —OH group... 

Nitrate ion is stabilized by electron delocalization which we can represent m terms of resonance between three equivalent Lewis structures... 

This chapter sets the stage for all of the others by reminding us that the relationship between structure and properties is what chemistry is all about It begins with a review of Lewis structures moves to a discussion of the Arrhenius Brpnsted-Lowry and Lewis pictures of acids and bases and the effects of structure on acidity and basicity... 

Section 1 3 The most common kind of bonding involving carbon is covalent bond ing A covalent bond is the sharing of a pair of electrons between two atoms Lewis structures are written on the basis of the octet rule, which limits second row elements to no more than eight electrons m their valence shells In most of its compounds carbon has four bonds... 

Table 14m this section sets forth the procedure to be followed m writ mg Lewis structures for organic molecules It begins with experimentally determined information the molecular formula and the constitution (order m which the atoms are connected)... 

Many molecules can be represented by two or more Lewis structures that differ only m the placement of electrons In such cases the electrons... 

Two Lewis structures (resonance forms) of formamide the atoms are connected in the same order but the arrangment of the electrons is different... 

Electron delocalization m the conjugate base usually expressed via resonance between Lewis structures increases acidity... 

All the following compounds are charactenzed by ionic bonding between a group I metal cation and a tetrahedral anion Wnte an appropriate Lewis structure for each anion remembenng to specify formal charges where they exist... 

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