# Lewis Structure Examples

### General | Latest Info

Neutral hydrogen atom has one valence electron. Each hydrogen atom in molecule shares one pair of bonding electrons and is therefore assigned one electron. Using Equation 8. 52 to calculate formal charge on Hydrogen, We obtain calculate formal charges on each atom of NH 4 + ion. Identify the number of valence electrons in each atom in NH 4 + ion. Use the Lewis electron structure of NH 4 + to identify the number of bonding and nonbonding electrons associated with each atom and then use Equation 8. 52 to calculate the formal charge on each atom. The Lewis electron structure for NH 4 + ion is as follow: nitrogen atom shares four bonding pairs of electrons, and the neutral nitrogen atom has five valence electrons. Using Equation 8. 51, formal charge on nitrogen atom is therefore {matheq} formal\; charge\left ( N \right )=5-\left ( 0+\frac{8}{2} \right )=0 {endmatheq} Each hydrogen atom has one bonding pair. The formal charge on each hydrogen atom is therefore {matheq} formal\; charge\left ( H \right )=1-\left ( 0+\frac{2}{2} \right )=0 {endmatheq} formal charges on atoms in NH 4 + ion. Thus, adding together formal charges on atoms should give us total charge on molecule or ion. In this case, sum of formal charges is 0 + 1 + 0 + 0 + 0 = + 1. Thiocyanate ion, which is used in printing and as a corrosion inhibitor against acidic gases, has at least two possible Lewis electron structures. Draw two possible structures, assign formal charges on all atoms in both, and decide which is the preferred arrangement of electrons. Ask for: Lewis electron structures, formal charges, and preferred arrangement use step - by - step procedure to write two plausible Lewis electron structures for SCN. Calculate formal charge on each atom using Equation 8. 51 Predict which structure is preferred based on formal charge on each atom and its electronegativity relative to other atoms present. Possible Lewis structures for SCN ion are as follow: b We must calculate formal charges on each atom to identify a more stable structure. If we begin with carbon, we notice that the carbon atom in each of these structures shares four bonding pairs, number of bonds typical for carbon, SO it has a formal charge of zero. Continuing with sulfur, we observe that in sulfur atom shares one bonding pair and has three lone pairs and has a total of six valence electrons. The formal charge on sulfur atom is therefore {matheq} 6-\left ( 6+\frac{2}{2} \right )=-1.5-\left ( 4+\frac{4}{2} \right )=-1 {endmatheq} in, nitrogen has a formal charge of −2. C Which structure is prefer? Structure is preferred because negative charge is on more electronegative atom, and it has lower formal charges on each atom as compared to structure: 0 1 versus + 1 2.

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### Definition

#### Table

N5
O (x 3)18
charge1
24

Lewis Structure is based on the concept of the octet rule, in which atoms share electrons so that each atom has eight electrons in its outer shell. As example, oxygen Atom has six electrons in its outer shell. In the Lewis Structure, these six dots are arranged so that Atom has two lone pairs and two single electrons. Two pairs would be opposite each other around the O symbol and two single electrons would be on other sides of Atom, opposite each other. In general, single electrons are written on the side of the element symbol. Incorrect placement would be, four electrons on one side of the Atom and two on the opposite side. When oxygen bonds to two hydrogen atoms to form water, each hydrogen atom has one dot for its lone electron. The Electron dot structure for water shows single electrons for oxygen sharing space with single electrons from hydrogen. All eight spots for dots around oxygen are fill, so the molecule has a stable octet.

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### How It Works

#### Table

N5
O (x 3)18
charge1
24

Draw out the correct Lewis Structure for the following compounds. Hcn LiF C 3 H 6 CO 3 2 - CH 3 NO 2 Another simple and general procedure to draw Lewis Structures has been proposed by ABP Lever. Before beginning this procedure, it is necessary to know the basic geometry of a molecule, ie whether it is cyclic or noncyclic, and which atoms are connected to which. Several work examples for determination of Lewis Structures of simple and more complicated species using the above described method can be found in the original paper and in Chemistry Net.

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### How to Write One

While Lewis Structures are useful when you re learning about valence, oxidation states, and bondingthere, there are many exceptions to rules in the real world. Atoms seek to fill or half - fill their Valence electron shell. However, atoms can and do form molecules that are not ideally stable. In some cases, central atom can form more than other atoms connect to it. The number of Valence Electrons can exceed eight, especially for higher atomic numbers. Lewis Structures are helpful for light elements but less useful for transition metals such as lanthanides and actinides. Students are caution to remember Lewis Structures are valuable tool for learning about and predicting behavior of atoms in molecules, but they are imperfect representations of real electron activity.

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### Lewis Symbols

We use Lewis symbols to describe valence electron configurations of atoms and monatomic ions. Lewis symbols consist of an elemental symbol surrounded by one dot for each of its valence electrons: figure 1 shows Lewis symbols for elements of the third period of the periodic table. Lewis symbols can also be used to illustrate formation of cations from atoms, as shown here for sodium and calcium: likewise, they can be used to show formation of anions from atoms, as shown here for chlorine and sulfur: figure 2 demonstrates use of Lewis symbols to show transfer of electrons during formation of ionic compounds.

##### Lewis Structures

For very simple molecules and molecular ions, we can write Lewis structures by merely pairing up unpaired electrons on constituent atoms. See these examples: For more complicated molecules and molecular ions, it is helpful to follow the step - by - step procedure outlined here: determining total number of valence electrons among all atoms. For cations, subtract one electron for each positive charge. For anions, add one electron for each negative charge. Draw skeleton structure of a molecule or ion, arranging atoms around the central atom. Connect each atom to the central atom with a single bond. Distribute remaining electrons as lone pairs on terminal atoms, completing octet around each atom. Place all remaining electrons on the central atom. Rearrange electrons OF outer atoms to make multiple bonds with central atom in order to obtain octets wherever possible. Let us determine Lewis structures OF SiH 4, {matheq}{CHO2-}{endmatheq} NO +, and OF 2 as examples in following this procedure: determine total number OF valence electrons in molecule or ion. For molecule, we add the number OF valence electrons on each atom in molecule: {matheq}\begin{align} &\phantom{+}{SiH4}\ &\phantom{+}\textrm{Si: 4 valence electrons/atom × 1 atom = 4}\ &\underline{\textrm{+H: 1 valence electron/atom × 4 atoms = 4}}\ &\hspace{271px}\textrm{= 8 valence electrons} \end{align}{endmatheq} For negative ion, such as {matheq}{CHO2-}{endmatheq} we add the number OF valence electrons on atoms to the number of negative charges on ion: {matheq}{CHO2-}\ \textrm{C: 4 valence electrons/atom × 1 atom} \hspace{6px}= \phantom{1}4\ \textrm{H: 1 valence electron/atom × 1 atom} \hspace{12px}= \phantom{1}1\ \textrm{O: 6 valence electrons/atom × 2 atoms = 12}\ \underline{+\hspace{100px}\textrm{1 additional electron} \hspace{9px}= \phantom{1}1}\ \hspace{264px}\textrm{= 18 valence electrons}{endmatheq} For positive ion, such as NO +, we add the number OF valence electrons on atoms in ion and then subtract number OF positive charges on ion from total number OF valence electrons: {matheq}{NO+}\ \textrm{N: 5 valence electrons/atom × 1 atom} = \phantom{−}5\ \textrm{O: 6 valence electron/atom × 1 atom}\hspace{5px} = \phantom{−}6\ \underline{\textrm{+ −1 electron (positive charge)} \hspace{44px}= −1}\ \hspace{260px}\textrm{= 10 valence electrons}{endmatheq} Since OF 2 is neutral molecule, we simply add number OF valence electrons: {matheq}\phantom{+ }{OF2}\ \phantom{+ }\textrm{O: 6 valence electrons/atom × 1 atom} \hspace{10px}= 6\ \underline{\textrm{+ F: 7 valence electrons/atom × 2 atoms} = 14}\ \hspace{280px}\textrm{= 20 valence electrons}{endmatheq} Draw skeleton structure OF molecule or ion, arranging atoms around central atom and connecting each atom to central atom with single bond. When several arrangements OF atoms are possible, as for {matheq}{CHO2-}{endmatheq} we must use experimental evidence to choose the correct one. In general, less electronegative elements are more likely to be central atoms. In {matheq}{CHO2-}{endmatheq} less electronegative carbon atom occupies central position with oxygen and hydrogen atoms surrounding it. Other examples include P in POCl 3, S in SO 2, and Cl. In {matheq}{ClO4-}{endmatheq} exception is that hydrogen is almost never the central atom. Like most electronegative element,ss fluorine also cannot be central atom. Distribute remaining electrons as lone pairs on terminal atoms to complete their valence shells with octet OF electrons. There are NO remaining electrons on SiH 4, SO it is unchanged: Place all remaining electrons on the central atom. For SiH 4, {matheq}{CHO2-}{endmatheq} and NO +, there are NO remaining electrons; we already place all OF electrons determined in Step 1. For OF 2, we had 16 electrons remaining in Step 3, and we Place 12, leaving 4 to be Place on the central atom: Rearrange electrons OF outer atoms to make multiple bonds with the central atom in order to obtain octets wherever possible.

##### The Octet Rule

Other halogen molecules form bonds like those in chlorine molecule: one single bond between atoms and three lone pairs of electrons per atom. This allows each halogen atom to have a noble gas electron configuration. The tendency of main group atoms to form enough bonds to obtain eight valence electrons is known as the octet rule. The number of bonds that atom can form can often be predicted from the number of electrons needed to reach octet; this is especially true of nonmetals of second period of the periodic table. For example, each atom of group 14 elements has four electrons in its outermost shell and therefore requires four more electrons to reach the octet. These four electrons can be gained by forming four covalent bonds, as illustrated here for carbon in CCl 4 and silicon in SiH 4. Because hydrogen only needs two electrons to fill its valence shell, it is an exception to the octet rule and only needs to form one bond. Transition elements and inner transition elements also do not follow the octet rule since they have d and F electrons involved in their valence shells. Group 15 elements such as nitrogen have five valence electrons in atomic Lewis symbol: one lone pair and three unpaired electrons. To obtain octet, these atoms form three covalent bonds, as in NH 3. Oxygen and other atoms in group 16 obtain octets by forming two covalent bonds:

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### Key Concepts and Summary

We use Lewis symbols to describe valence electron configurations of atoms and monatomic ions. Lewis symbols consist of an elemental symbol surrounded by one dot for each of its valence electrons: Figure 1 shows Lewis symbols for elements of the third period of the periodic table. Lewis symbols can also be used to illustrate the formation of cations from atoms,. As shown here for sodium and calcium: Likewise, they can be used to show formation of anions from atoms,. As shown below for chlorine and sulfur: Figure 2 demonstrates use of Lewis symbols to show transfer of electrons during formation of ionic compounds.

##### Lewis Structures

Nasa's Cassini - Huygens mission detected a large cloud of toxic hydrogen cyanide on Titan, one of Saturn's moons. Titan also contains ethane, acetylene, and ammonia. What are Lewis structures of these molecules? Step 1: Calculate the number of valence electrons. Hcn: + = 10 H 3 CCH 3: + = 14 HCCH: + = 10 NH 3: + = 8 step 2. Draw skeleton and connect atoms with single bonds. Remember that H is never a central atom: step 3: Where needed to distribute electrons to terminal atoms: step 4: Where needed to place remaining electrons on the central atom:

##### The Octet Rule

Other halogen molecules form bonds like those in chlorine molecule: one single bond between atoms and three lone pairs of electrons per atom. This allows each halogen atom to have a noble gas electron configuration. The tendency of main group atoms to form enough bonds to obtain eight valence electrons is known as the octet rule. The number of bonds that atom can form can often be predicted from the number of electrons needed to reach octet; this is especially true of nonmetals of second period of the periodic table. For example, each atom of group 14 elements has four electrons in its outermost shell and therefore requires four more electrons to reach the octet. These four electrons can be gained by forming four covalent bonds, as illustrated here for carbon in CCl 4 and silicon in SiH 4. Because hydrogen only needs two electrons to fill its valence shell, it is an exception to the octet rule. Transition elements and inner transition elements also do not follow the octet rule: group 15 elements such as nitrogen have five valence electrons in atomic Lewis symbol: one lone pair and three unpaired electrons. To obtain octet, these atoms form three covalent bonds, as in NH 3. Oxygen and other atoms in group 16 obtain octets by forming two covalent bonds:

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### Lewis Structures for Polyatomic Ions

We also use Lewis symbols to indicate the formation of covalent bonds, which are shown in Lewis structures, drawings that describe bonding in molecules and polyatomic ions. For example, when two chlorine atoms form chlorine molecule, they share one pair of electrons: Lewis structure indicates that each Cl atom has three pairs of electrons that are not used in bonding and one share pair of electrons. Dash is sometimes used to indicate shared pair of electrons: single shared pair of electrons is called a single bond. Each Cl atom interacts with eight valence electrons: six in lone pairs and two in single bond.

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### What is Electron Dot Structure?

#### Table

BondBond Length
N-N1.47 A
N=N1.24 A
NN1.10 A

The Lewis structure is a structural representation of molecule where dots are used to show electron positions around atoms and lines or dot pairs represent covalent bonds between atoms. The purpose of drawing Lewis dot structure is to identify lone electron pairs in molecules to help determine chemical bond formation. Lewis structures can be made for molecules that contain covalent bonds and for coordination compounds. The reason is that electrons are shared in covalent bond. In ionic bond, it's more like one atom donates an electron to other atom. Lewis structures are named for Gilbert N. Lewis, who introduced the idea in the article Atom and Molecule in 1916. Also known as: Lewis structures are also called Lewis dot diagrams, electron dot diagrams, Lewis dot formulas, or electron dot formulas. Technically, Lewis structures and electron dot structures are different because electron dot structure show all electrons as dots, while Lewis structures indicate share pairs in chemical bond by drawing line.

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### The Octet Rule

Why are some substances chemically bond molecules and others are association with ions? The answer to this question depends upon electronic structures OF atoms and the nature of chemical forces within compounds. Although there are NO sharply defined boundaries, chemical bonds are typically classified into three main types: ionic bonds, Covalent bonds, and Metallic bonds. In this chapter, each type of bond will be discussed and general properties found in typical substances in which bond type occur ionic bonds result from electrostatic forces that exist between ions of opposite charge. These bonds typically involve metal with nonmetal - covalent bonds that result from sharing OF electrons between two atoms. Bonds typically involve one nonmetallic element with another metallic bond. These bonds are found in solid metal with each metal bond to several neighboring groups and bonding electrons free to move throughout the 3 - dimensional structure. Each bond classification is discussed in detail in subsequent sections OF chapter. Let's look at preferred arrangements of electrons in atoms when they form chemical compounds. Figure 8. 11: G. N. Lewis and Octet Rule. Lewis is working in a laboratory. In Lewis ' original sketch For Octet Rule, he initially placed electrons at corners of the cube rather than placing them as we do now. In 1904, Richard Abegg formulated what is now known as Abegg's Rule, which states that the difference between maximum positive and negative valences of element is frequently eight. This rule was used later in 1916 when Gilbert N. Lewis formulated the Octet Rule in his cubical atom theory. The Octet Rule refers to the tendency OF atoms to prefer to have eight electrons in valence shell. When atoms have fewer than eight electrons, they tend to react and form more stable compounds. Atoms will react to get in most stable state possible. Complete Octet is very stable because all orbitals will be full. Atoms with greater stability have less energy, so reactions that increase stability of atoms will release energy in the form of heat or light; reactions that decrease stability must absorb energy, getting colder. When discussing the Octet Rule, we do not consider d or F electrons. Only S and P electrons are involved in the Octet Rule, making it a useful rule for main Group elements; Octet in these atoms corresponds to electron configurations ending with S 2 P 6. Lewis dot symbols can also be used to represent ions in ionic compounds. Reaction of cesium with fluorine, For example, to produce ionic compound CsF can be written as follow: NO dots are shown on Cs + in product because cesium has lost its single valence electron to fluorine. Transfer OF this electron produces Cs + ion, which has a valence electron configuration OF Xe, and F ion, which has a total OF eight valence electrons and Ne electron configuration.

##### Electron-deficient Molecules

We will also encounter a few molecules that contain central atoms that do not have fill valence shell. Generally, these are molecules with central atoms from groups 2 and 12, outer atoms that are hydrogen, or other atoms that do not form multiple bonds. For example, in Lewis structures of beryllium dihydride, BeH 2, and boron trifluoride, BF 3, beryllium and boron atoms each have only four and six electrons, respectively. It is possible to draw a structure with a double bond between boron atom and fluorine atom in BF 3, satisfying the octet rule, but experimental evidence indicates bond lengths are closer to that expected for B - F single bonds. This suggests the best Lewis structure has three B - F single bonds and electron deficient boron. Reactivity of compound is also consistent with electron deficient boron. However, B - F bonds are slightly shorter than what is actually expected for B - F single bonds, indicating that some double bond characters are found in actual molecule. Atoms like boron atom in BF 3, which do not have eight electrons, are very reactive. It readily combines with molecule containing atom with a lone pair of electrons. For example, NH 3 reacts with BF 3 because lone pair of nitrogen can be shared with boron atom:

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### Double and Triple Bonds

Single Bond: Covalent Bond between two atoms that is formed by one bonding pair of electrons. G - 14 double Bond: multiple Covalent Bond between two atoms that is formed by two bonding electron pairs. G - 5 triple Bond: multiple Covalent Bond between two atoms that is formed by three bonding pairs of electrons. G - 15 multiple Covalent Bond: more than one Bond between atoms, formed by sharing of more than one electron from each atom. It can be either double Bond or triple Bond. G - 10 resonance Structures: two or more equally correct dot representations of the same molecule; representation is the actual structure of the molecule.

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### Example 1

Periodic table has all of information needed to Draw Lewis Dot Structure. Each group, or column, is indicated by a roman numeral which represents the number of Valence Electrons. This is applicable to entire group. For example, all elements which fall within the first column, or Group I, have one Valence electron. All elements in Group II have two Valence Electrons, all way up to VIII, eight Valence Electrons. Properties are also consistent across rows, or periods, of periodic table. Periods are indicated by number, 1 2 3, etc. Which represents the energy level, or shell of electrons. The first period, or row, has only one energy level that can hold a total of two electrons. Period 2, with a second shell, can hold a total of eight electrons, also know as the octet rule. Period 3 and so forth can hold more than eight electrons. Periodic tables also convey electronegativity. Most electronegative elements are located in uppermost right corner of the period table and decrease in electronegativity as you go down Group or more left of period. Throughout drawing Lewis Dot Structures, periodic table will be a strong reference point when working with Electrons, covalent bonding, and polyatomic ions. Assume that each outer element has full valence of bonding and non - bonding electrons. Total all of these electrons, and subtract that from the total number of Valence Electrons in the molecule. Co 2 has 16 Valence Electrons. We assume each O has 8 Valence Electrons. 28 = 16; 16 - 16 = 0. Therefore, we do need to add any electrons to C.

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