Lewis Model of Electronic Structure · 2/8/2013 · Activity 08-2 Lewis Model of Electronic Structure 147 EXERCISES 5. Draw the Lewis structures for O 2, N 2, C 2 H 4, and C 2 H

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WHY? The Lewis model of molecular electronic structure describes how atoms bond to each other to form molecules. It determines the number of bonds formed between pairs of atoms in a molecule and the number of electrons that exist as lone or nonbonding pairs. This information makes it possible for you to predict the geometry of molecules (e.g., CO2 is linear but SO2 and H2O are bent) and relative bond strengths and lengths.

LEARNING OBJECTIVE• Understand how to draw Lewis structures and interpret them in terms of molecular properties

SUCCESS CRITERIA• Construct realistic Lewis structures• Identify relative bond strengths and lengths from Lewis structures

PREREQUISITES• Activity 03-1: Molecular Representations• Activity 07-4: Multi-electron Atoms, the Aufbau Principle, and the Periodic Table

INFORMATIONMolecules exist because they are more stable than separated atoms. By “more stable,” we mean that they have lower energy in the same way that a skateboarder at the bottom of a hill has less energy or is more stable than one at the top. The physical chemist G. N. Lewis recognized, from the very low chemical reactivity of the noble gases, that a confi guration with eight electrons in a shell produces a very stable situation. He therefore proposed that molecules form so that atoms can transfer or share electrons and produce this very stable octet structure.Lewis structures are used to model how the electrons are arranged in order to produce these stable eight-electron confi gurations. In these diagrams, dots are used to represent electrons; a line between atoms represents a single covalent bond formed by a pair of electrons; other dots represent nonbonding electrons; and charges are written to identify a formal distribution of charge.The bonds show how the atoms in a molecule are connected to each other. A Lewis diagram does not show bond lengths, bond angles, the arrangement of atoms in three-dimensional space, or the actual charges on atoms. Some molecules require more than one Lewis structure to describe them. These multiple structures are called resonance structures. In some situations, atoms in Period 2 can have fewer than 8 electrons and atoms in Periods 3 and higher have more than 8 electrons.Some atoms (e.g., C, N, O, and S) form double bonds, which are represented by double lines. Some atoms (e.g., C and N) form triple bonds, which are represented by triple lines. A double bond is stronger and shorter than single bond, and a triple bond is the strongest and shortest of the three.How does one determine and draw a Lewis structure? First determine whether the molecule is ionic or covalent. If it is ionic, draw each ion separately. For covalent molecules and polyatomic ions, follow the methodology given in Model 1.

Lewis Model of Electronic StructureACTIVITY 08-2

Chapter 8: Covalent Bonding144

Foundations of Chemistry

MODEL 1: METHODOLOGY FOR CONSTRUCTING LEWIS STRUCTURES HCl – A Simple Example

Methodology Example: Single Bond

Step 1: Count the valence electrons from all the atoms. Add electrons for negative ions, subtract electrons for positive ions. The number of valence electrons can be determined from the atom’s position in the Periodic Table.

For hydrochloric acid, H has 1 and Cl has 7 for a total of 8

valence electrons.

Step 2: Assemble the bonding framework. Decide which atoms are connected to each other, and use a pair of electrons, represented by a line, to form a bond between each pair of atoms that are bonded together.

H Cl

Step 3: Arrange the remaining electrons so that each atom has 8 electrons around it (the octet rule). If necessary, place additional pairs of electrons between the atoms to form additional bonds.

H Cl

Step 4: Check for exceptions to the octet rule. For H, only 2 electrons are needed (the duet rule). Be and B are often electron defi cient and may have only 4 or 6 electrons. Atoms in the second period cannot have more than 8 electrons. The third and higher period elements sometimes have more than 8 electrons.

Cl satisfi es the octet rule.H satisfi es the duet rule.

Step 5: Determine the formal charges (FC) on the atoms.FC = number of atomic valence electrons – number of lone pair electrons – 0.5(number of shared electrons).Evaluate whether the formal charges (FC) on the atoms are reasonable. The structure is reasonable if the charges are zero or small, and the negative charges reside on the most electronegative atoms. Other structures with larger formal charges are higher energy and are not representative of the lowest energy structure of the molecule.

FC(H) = 1 – 0 – 0.5 (2) = 0FC(Cl) = 7 – 6 – 0.5 (2) = 0

Reasonable becauseFC is zero.

Step 6: Draw resonance structures. none in this case

KEY QUESTIONS1. How can you determine the number of valence electrons that an atom has?

Activity 08-2 Lewis Model of Electronic Structure 145

2. a) How many valence electrons does H have and how many more does it need to fi ll the fi rst shell?

b) How many valence electrons does Cl have and how many more does it need to achieve a noble gas confi guration?

3. From the Lewis perspective, why do hydrogen and chlorine combine to form the HCl molecule?

4. How is formal charge determined?

5. How is formal charge used to identify unreasonable Lewis structures?

EXERCISES1. Write the number of valence electrons for H and for F. Draw the Lewis structure for HF.

2. Write the number of valence electrons for C. Determine how many more electrons C needs to form

a molecule, and draw the Lewis structure for a molecule composed of C and H.

3. Write the number of valence electrons for S. Determine how many more electrons S needs to form a molecule, and draw the Lewis structure for a molecule composed of S and H.



4. Draw Lewis structures for H2O, NH3, PCl3, C2H6, and NaCl. (NaCl is ionic, so draw the two ions separately, with no line connecting them, and indicate the charge on the ion with a + or − sign.)

Table 1CO2 – Need for Double Bonds

Methodology Example: Multiple Bonds

Step 1: Count the number of valence electrons from all the atoms. For carbon dioxide, C has 4 valence electrons and O has 6 valence electrons for a total

of 16 valence electrons.

Step 2: Assemble the bonding framework. C OOStep 3: Arrange the remaining electrons so that each atom has

8 electrons around it (the octet rule). If necessary, place additional pairs of electrons between the atoms to form additional bonds to satisfy the octet rule.

C OO

Step 4: Check for exceptions to the octet rule. No exceptions present.

Step 5: Evaluate whether the formal charges on the atoms are reasonable.

FC(C) = 4 – 0 – 0.5 (8) = 0FC(O) = 6 – 4 – 0.5 (4) = 0

Reasonable becauseFC is zero.


KEY QUESTIONS 6. How many valence electrons are there in the carbon dioxide molecule?

7. Why is it sometimes necessary in constructing Lewis structures to put double or even triple bonds between atoms?


EXERCISES5. Draw the Lewis structures for O2, N2, C2H4, and C2H2.

ICl4– – Exception to the Octet Rule Table 2

Methodology Example:Third Period Element

Step 1: Count the number of valence electrons from all the atoms, and, because it is a –1 anion, add one.

For ICl4–, there are 36 valence electrons (5 × 7 + 1)

Step 2: Assemble the bonding framework. ICl

Cl

Cl

Cl

Step 3: Arrange the remaining electrons so that each atom has 8 electrons around it. I ..

.. :Cl

..

.. :Cl

..

..:Cl

..

..:Cl

Step 4: Check for exceptions to the octet rule. After satisfying the octet rule for each atom, 4 electrons remain. These are placed as nonbonding electrons on the fi fth period element, iodine.

..I.. ..

.. :Cl

..

.. :Cl

..

..:Cl

..

..:Cl-


Note: The negative formal charge is on I, which is less electronegative than Cl. This result is correct and is an exception to the general rule given earlier.

FC(I) = 7 – 8 = –1FC(Cl) = 7 – 7 = 0




KEY QUESTIONS8. How does the example of ICl4– differ from the examples of HCl and CO2?

EXERCISE6. Draw the Lewis structure for PCl5.

NO2– – The Need for Multiple Lewis Structures Table 3

Methodology Example: Resonance

Step 1: Count the number of valence electrons from all the atoms, and, because it is a –1 ion, add one.

For NO2– there are 18 valence electrons (5 + (2 × 6) + 1)

Step 2: Assemble the bonding framework. N OO

Step 3: Arrange the remaining electrons so that each atom has 8 electrons around it. N OO

-

Step 4: Check for exceptions to the octet rule. none


FC(N) = 5 – 5 = 0FC(O1) = 6 – 7 = –1FC(O2) = 6 – 6 = 0

Step 6: In Step 3 the double bond could just as well have been drawn on the left side of the nitrogen. This second structure is called a resonance structure. In fact, experimental measurements show that both N-O bonds are the same with a bond energy and a bond length that are intermediate between a typical N-O single bond and a typical N=O double bond. Therefore, the second structure is needed to adequately describe the bonding. Draw the resonance structure to show that both bonds are equivalent.

N OO-


KEY QUESTIONS9. Why is it not possible to describe NO2– by a single Lewis structure?

EXERCISE7. Draw the Lewis structure for ozone O3. Include the formal charge in parentheses above each atom

if the formal charge differs from zero, e.g., (+1) or (–1).

MODEL 2: BOND STRENGTHS AND LENGTHS

Bond Average Dissociation Energy in kJ/mole Average Length in pm

C−C 345 154

C=C 615 133

C≡C 835 120

KEY QUESTIONS10. The energy it takes to dissociate or break a bond is a measure of the bond strength. In view of the

data in Model 2, how does the Lewis structure help you identify the strongest bonds in a molecule?

11. In view of the data in Model 2, how does the Lewis structure help you identify the shortest bonds in a molecule?



EXERCISE8. Identify which C–O bond in acetic acid is the shortest and strongest.

O

H O

H

HCH C

MODEL 3: WHICH LEWIS STRUCTURE IS BETTER?The two structures below represent the thiocyanate ion. Which one has a lower energy and is therefore the better description of the lowest energy state of this ion?

C NS

(A)

S C N

(B)

KEY QUESTIONS12. What are the formal charges on the atoms in the thiocyanate ion in Model 3? Write the formal

charge in parentheses above each atom in these structures.

13. Based on the criteria developed in Key Question 5, which structure for thiocyanate (in Model 3)has the lower energy and is, therefore, the better description of the lowest energy state of this ion?


EXERCISES9. In view of the formal charges, why is O = C = O a better representation of carbon dioxide than

O − C ≡ O ?

INFORMATIONIn assembling the bonding framework of a molecule, you may fi nd it diffi cult to identify which atoms are bonded to each other. Sometimes this information cannot be deduced from fi rst principles: you simply need to know the molecular structure. Here are some guidelines that are often helpful.

• It is useful to think in terms of outer atoms and inner atoms. An outer atom bonds to only one other atom, while an inner atom bonds to more than one other atom.

• Hydrogen atoms are always outer atoms because they can form only one bond.

• Outer atoms other than hydrogen are usually the ones with the highest electronegativities.

• The bonding framework is often indicated by the order in which the atoms are written in a molecular formula. For example, in OCS the carbon atom is the inner atom.

• Parentheses are often used in a molecular formula to indicate the bonding framework. For example in (CH3)2CO, the hydrogen atoms are bonded to carbon to form two methyl radicals, and oxygen is an outer atom bonded to an inner carbon atom.

• Multiple atoms of the same element are usually the outer atoms around a single atom of another element. For example, in PF6 , phosphorus is the inner atom.

• Sometimes it is helpful to know the chemical properties of a molecular formula. For example, in HNO3 the hydrogen atom might be bonded to the nitrogen or to oxygen. If you know that nitric acid is an oxyacid, you can locate the hydrogen as bonded to an oxygen atom.

• Finally, the most likely structure is the one which has the most reasonable formal charges on the atoms. By “reasonable,” we mean that the formal charges should be small or zero, and the negative charges should be located on the most electronegative atoms.



EXERCISES10. Draw Lewis structures for the following molecules. A Lewis structure includes the formal charge on

each atom if it differs from zero and any resonance structures that are signifi cant. To be signifi cant, a resonance structure does not increase the formal charge on any of the atoms. Indicate the formal charge on an atom by writing it in parentheses next to the atom, e.g., F (–1).

SeF2 I3– OCS

HNO3 SiCl4

11. Use Lewis structures to arrange the following compounds in order of increasing carbon-carbon bond strength. Explain.

C2H6 (ethane) C2H4 (ethylene) C2H2 (acetylene)

12. Use Lewis structures to arrange the following compounds in order of increasing carbon-carbon bond length. Explain.

C2H6 (ethane) C2H4 (ethylene) C2H2 (acetylene)


13. Explain which N–N bond is more stable: the one in nitrogen (N2) or the one in hydrazine (N2H4).

14. Explain which C–O bond is shorter: the one in methanol (CH3OH) or the one in formaldehyde (CH2O).

PROBLEMS1. Use Lewis structures to identify which of the following compounds you would most likely be

successful at synthesizing: SiF4, OF4, SF6, or OF6. Explain.

2. Two Lewis structures are needed to describe the bonding in formamide, HCONH2. Write these two resonance structures. One should have no formal charges on the atoms, and the other will have a formal charge of +1 on N and –1 on O. Of the structures you have drawn, which do you expect to be the better description of formamide?



3. For the two resonance structures of formamide in Problem 2, explain why each of the following statements is either correct or incorrect.a) The molecule is not oscillating back-and-forth between the two structures; the molecule is,

instead, an average or superposition of the two structures.

b) The expected CO bond length is between that for a normal CO double bond and that for a normal CO single bond.

c) The nitrogen in the molecule has a lower electron density associated with it than is found for the free nitrogen atom.

d) Both resonance structures have the same energy.

REFLECTIONDevelop a checklist that you can use to ensure that you have written a correct Lewis structure.

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Lewis Model of Electronic Structure · 2/8/2013 · Activity 08-2 Lewis Model of Electronic Structure 147 EXERCISES 5. Draw the Lewis structures for O 2, N 2, C 2 H 4, and C 2 H