how do you draw a single bond in a lewis structure?

Affiliate seven. Chemical Bonding and Molecular Geometry

vii.three Lewis Symbols and Structures

Learning Objectives

By the end of this section, yous will be able to:

• Write Lewis symbols for neutral atoms and ions
• Depict Lewis structures depicting the bonding in uncomplicated molecules

Thus far in this chapter, nosotros have discussed the various types of bonds that form between atoms and/or ions. In all cases, these bonds involve the sharing or transfer of valence vanquish electrons between atoms. In this section, nosotros volition explore the typical method for depicting valence shell electrons and chemical bonds, namely Lewis symbols and Lewis structures.

Lewis Symbols

Nosotros utilize Lewis symbols to draw valence electron configurations of atoms and monatomic ions. A Lewis symbol consists of an elemental symbol surrounded by 1 dot for each of its valence electrons:

Effigy 1 shows the Lewis symbols for the elements of the third period of the periodic table.

Figure i. Lewis symbols illustrating the number of valence electrons for each chemical element in the 3rd period of the periodic table.

Lewis symbols can likewise be used to illustrate the germination of cations from atoms, as shown here for sodium and calcium:

As well, they can exist used to evidence the germination of anions from atoms, as shown here for chlorine and sulfur:

Figure 2 demonstrates the use of Lewis symbols to evidence the transfer of electrons during the formation of ionic compounds.

Figure 2. Cations are formed when atoms lose electrons, represented by fewer Lewis dots, whereas anions are formed past atoms gaining electrons. The full number of electrons does non change.

Lewis Structures

We also apply Lewis symbols to betoken the formation of covalent bonds, which are shown in Lewis structures, drawings that describe the bonding in molecules and polyatomic ions. For example, when two chlorine atoms form a chlorine molecule, they share one pair of electrons:

The Lewis structure indicates that each Cl atom has 3 pairs of electrons that are not used in bonding (chosen lone pairs) and one shared pair of electrons (written between the atoms). A dash (or line) is sometimes used to indicate a shared pair of electrons:

A unmarried shared pair of electrons is chosen a single bail. Each Cl atom interacts with eight valence electrons: the half-dozen in the lone pairs and the two in the single bond.

The Octet Dominion

The other halogen molecules (F_ii, Br₂, I₂, and At_two) class bonds similar those in the chlorine molecule: 1 single bond between atoms and 3 alone pairs of electrons per atom. This allows each halogen atom to have a element of group 0 electron configuration. The tendency of primary grouping atoms to form enough bonds to obtain eight valence electrons is known as the octet rule.

The number of bonds that an cantlet can form can often be predicted from the number of electrons needed to reach an octet (8 valence electrons); this is peculiarly true of the nonmetals of the second period of the periodic table (C, N, O, and F). For example, each atom of a group 14 chemical element has four electrons in its outermost shell and therefore requires 4 more than electrons to achieve an octet. These four electrons tin can be gained by forming four covalent bonds, every bit illustrated here for carbon in CCl₄ (carbon tetrachloride) and silicon in SiH_iv (silane). Considering hydrogen simply needs two electrons to fill its valence vanquish, information technology is an exception to the octet rule. The transition elements and inner transition elements as well do not follow the octet rule:

Grouping 15 elements such every bit nitrogen have five valence electrons in the diminutive Lewis symbol: one lone pair and three unpaired electrons. To obtain an octet, these atoms class 3 covalent bonds, equally in NH₃ (ammonia). Oxygen and other atoms in group xvi obtain an octet past forming 2 covalent bonds:

Double and Triple Bonds

As previously mentioned, when a pair of atoms shares i pair of electrons, nosotros phone call this a unmarried bond. Still, a pair of atoms may need to share more than one pair of electrons in society to attain the requisite octet. A double bond forms when two pairs of electrons are shared betwixt a pair of atoms, equally between the carbon and oxygen atoms in CH₂O (formaldehyde) and between the two carbon atoms in C₂H₄ (ethylene):

A triple bond forms when three electron pairs are shared by a pair of atoms, as in carbon monoxide (CO) and the cyanide ion (CN^–):

Writing Lewis Structures with the Octet Rule

For very simple molecules and molecular ions, we tin write the Lewis structures by merely pairing up the unpaired electrons on the elective atoms. Come across these examples:

For more complicated molecules and molecular ions, information technology is helpful to follow the pace-by-step procedure outlined here:

1. Determine the total number of valence (outer beat) electrons. For cations, subtract ane electron for each positive charge. For anions, add together one electron for each negative charge.
2. Depict a skeleton structure of the molecule or ion, arranging the atoms around a central atom. (Generally, the least electronegative element should be placed in the center.) Connect each cantlet to the primal cantlet with a single bond (i electron pair).
3. Distribute the remaining electrons equally solitary pairs on the terminal atoms (except hydrogen), completing an octet around each cantlet.
4. Place all remaining electrons on the primal atom.
5. Rearrange the electrons of the outer atoms to brand multiple bonds with the central cantlet in order to obtain octets wherever possible.

Permit us decide the Lewis structures of SiH₄, CHO₂−, NO⁺, and OF₂ as examples in following this procedure:

1. Determine the total number of valence (outer shell) electrons in the molecule or ion.
   • For a molecule, we add the number of valence electrons on each atom in the molecule:

     [latex]\begin{assortment}{r r l} \text{SiH}_4 & & \\[1em] & \text{Si: iv valence electrons/atom} \times 1 \;\text{atom} & = 4 \\[1em] \dominion[-0.5ex]{21em}{0.1ex}\hspace{-21em} + & \text{H: ane valence electron/atom} \times 4 \;\text{atoms} & = 4 \\[1em] & & = 8 \;\text{valence electrons} \end{array}[/latex]

   • For a negative ion, such as CHO₂ ⁻, we add the number of valence electrons on the atoms to the number of negative charges on the ion (one electron is gained for each unmarried negative charge):

     [latex]\begin{array}{r r l} {\text{CHO}_2}^{-} & & \\[1em] & \text{C: 4 valence electrons/atom} \times 1 \;\text{atom} & = 4 \\[1em] & \text{H: 1 valence electron/atom} \times ane \;\text{atom} & = 1 \\[1em] & \text{O: 6 valence electrons/atom} \times 2 \;\text{atoms} & = 12 \\[1em] \rule[-0.5ex]{21.5em}{0.1ex}\hspace{-21.5em} + & 1\;\text{additional electron} & = ane \\[1em] & & = xviii \;\text{valence electrons} \end{array}[/latex]

   • For a positive ion, such as NO⁺, we add the number of valence electrons on the atoms in the ion and and so decrease the number of positive charges on the ion (one electron is lost for each unmarried positive accuse) from the full number of valence electrons:

     [latex]\begin{array}{r r l} \text{NO}^{+} & & \\[1em] & \text{Northward: 5 valence electrons/atom} \times one \;\text{atom} & = v \\[1em] & \text{O: half-dozen valence electrons/atom} \times 1 \;\text{cantlet} & = half dozen \\[1em] \rule[-0.5ex]{21em}{0.1ex}\hspace{-21em} + & -1 \;\text{electron (positive accuse)} & = -one \\[1em] & & = 10 \;\text{valence electrons} \cease{array}[/latex]

   • Since OF₂ is a neutral molecule, we simply add the number of valence electrons:

     [latex]\begin{array}{r r l} \text{OF}_{2} & & \\[1em] & \text{O: 6 valence electrons/cantlet} \times i \;\text{atom} & = 6 \\[1em] \dominion[-0.5ex]{21em}{0.1ex}\hspace{-21em} + & \text{F: 7 valence electrons/atom} \times 2 \;\text{atoms} & = xiv \\[1em] & & = 20 \;\text{valence electrons} \end{array}[/latex]

2. Draw a skeleton structure of the molecule or ion, arranging the atoms effectually a fundamental atom and connecting each atom to the cardinal cantlet with a unmarried (one electron pair) bail. (Notation that we denote ions with brackets around the structure, indicating the charge exterior the brackets:)When several arrangements of atoms are possible, every bit for CHO_two ⁻, we must employ experimental evidence to choose the correct ane. In general, the less electronegative elements are more likely to be primal atoms. In CHO₂ ⁻, the less electronegative carbon atom occupies the central position with the oxygen and hydrogen atoms surrounding it. Other examples include P in POCl₃, S in And so₂, and Cl in ClO₄ ⁻. An exception is that hydrogen is almost never a central cantlet. As the nigh electronegative element, fluorine also cannot be a central atom.
3. Distribute the remaining electrons every bit lone pairs on the concluding atoms (except hydrogen) to complete their valence shells with an octet of electrons.
   • There are no remaining electrons on SiH_iv, then it is unchanged:
4. Place all remaining electrons on the cardinal atom.
   • For SiH_iv, CHO₂ ⁻, and NO⁺, there are no remaining electrons; we already placed all of the electrons determined in Step 1.
   • For OF_two, we had sixteen electrons remaining in Pace 3, and we placed 12, leaving 4 to exist placed on the central atom:
5. Rearrange the electrons of the outer atoms to brand multiple bonds with the central cantlet in guild to obtain octets wherever possible.

Example i

Writing Lewis Structures
NASA'southward Cassini-Huygens mission detected a large cloud of toxic hydrogen cyanide (HCN) on Titan, one of Saturn'southward moons. Titan also contains ethane (H_threeCCH₃), acetylene (HCCH), and ammonia (NH₃). What are the Lewis structures of these molecules?

Solution

1. Summate the number of valence electrons.HCN: (1 × 1) + (four × 1) + (v × one) = 10H₃CCH_three: (1 × 3) + (ii × 4) + (1 × iii) = 14HCCH: (ane × i) + (ii × 4) + (1 × 1) = 10NH_three: (5 × 1) + (3 × 1) = 8
2. Draw a skeleton and connect the atoms with single bonds. Remember that H is never a central atom:
3. Where needed, distribute electrons to the terminal atoms: HCN: half-dozen electrons placed on NH₃CCH₃: no electrons remainHCCH: no final atoms capable of accepting electrons

   NH₃: no terminal atoms capable of accepting electrons

4. Where needed, place remaining electrons on the key atom: HCN: no electrons remainH₃CCH_three: no electrons remainHCCH: four electrons placed on carbon

   NH₃: ii electrons placed on nitrogen

5. Where needed, rearrange electrons to class multiple bonds in lodge to obtain an octet on each cantlet:HCN: grade two more C–N bondsH₃CCH₃: all atoms have the correct number of electronsHCCH: form a triple bond between the two carbon atomsNH₃: all atoms have the correct number of electrons

   

Cheque Your Learning
Both carbon monoxide, CO, and carbon dioxide, CO₂, are products of the combustion of fossil fuels. Both of these gases also cause problems: CO is toxic and CO_two has been implicated in global climatic change. What are the Lewis structures of these 2 molecules?

Answer:

Fullerene Chemistry

Carbon soot has been known to man since prehistoric times, just it was non until fairly recently that the molecular structure of the main component of soot was discovered. In 1996, the Nobel Prize in Chemistry was awarded to Richard Smalley (Figure 3), Robert Curl, and Harold Kroto for their work in discovering a new form of carbon, the C₆₀ buckminsterfullerene molecule (Figure 1 in Affiliate 7 Introduction). An entire class of compounds, including spheres and tubes of various shapes, were discovered based on C_60. This type of molecule, called a fullerene, shows promise in a variety of applications. Because of their size and shape, fullerenes can encapsulate other molecules, then they accept shown potential in various applications from hydrogen storage to targeted drug delivery systems. They likewise possess unique electronic and optical properties that accept been put to good apply in solar powered devices and chemic sensors.

Figure 3. Richard Smalley (1943–2005), a professor of physics, chemical science, and astronomy at Rice Academy, was one of the leading advocates for fullerene chemical science. Upon his death in 2005, the United states Senate honored him as the "Father of Nanotechnology." (credit: United States Department of Free energy)

Exceptions to the Octet Rule

Many covalent molecules have central atoms that do not have eight electrons in their Lewis structures. These molecules autumn into three categories:

• Odd-electron molecules have an odd number of valence electrons, and therefore accept an unpaired electron.
• Electron-deficient molecules have a central atom that has fewer electrons than needed for a noble gas configuration.
• Hypervalent molecules have a cardinal atom that has more electrons than needed for a noble gas configuration.

Odd-electron Molecules

We phone call molecules that comprise an odd number of electrons gratuitous radicals. Nitric oxide, NO, is an example of an odd-electron molecule; it is produced in internal combustion engines when oxygen and nitrogen react at loftier temperatures.

To depict the Lewis structure for an odd-electron molecule like NO, we follow the same five steps we would for other molecules, but with a few minor changes:

1. Determine the total number of valence (outer shell) electrons. The sum of the valence electrons is 5 (from North) + 6 (from O) = 11. The odd number immediately tells the states that nosotros take a free radical, and then nosotros know that not every atom tin accept eight electrons in its valence shell.
2. Draw a skeleton structure of the molecule. We can easily draw a skeleton with an Northward–O unmarried bond:Northward–O
3. Distribute the remaining electrons equally solitary pairs on the concluding atoms. In this case, there is no central atom, and then we distribute the electrons effectually both atoms. Nosotros give eight electrons to the more electronegative atom in these situations; thus oxygen has the filled valence crush:
   
4. Place all remaining electrons on the central atom. Since there are no remaining electrons, this step does not employ.
5. Rearrange the electrons to brand multiple bonds with the fundamental atom in order to obtain octets wherever possible. We know that an odd-electron molecule cannot take an octet for every cantlet, but we desire to get each atom as shut to an octet as possible. In this case, nitrogen has but 5 electrons around it. To move closer to an octet for nitrogen, we accept one of the lone pairs from oxygen and use information technology to form a NO double bond. (We cannot take another alone pair of electrons on oxygen and grade a triple bond because nitrogen would and so have ix electrons:)
   

Electron-scarce Molecules

Nosotros will also encounter a few molecules that contain central atoms that practise not have a filled valence shell. Generally, these are molecules with central atoms from groups 2 and 12, outer atoms that are hydrogen, or other atoms that practice not class multiple bonds. For example, in the Lewis structures of beryllium dihydride, BeH₂, and boron trifluoride, BF_three, the beryllium and boron atoms each have only four and half dozen electrons, respectively. It is possible to draw a structure with a double bond between a boron atom and a fluorine atom in BF₃, satisfying the octet rule, just experimental evidence indicates the bond lengths are closer to that expected for B–F single bonds. This suggests the best Lewis structure has 3 B–F single bonds and an electron deficient boron. The reactivity of the compound is also consistent with an electron deficient boron. Nevertheless, the B–F bonds are slightly shorter than what is actually expected for B–F single bonds, indicating that some double bond character is constitute in the actual molecule.

An atom like the boron atom in BF₃, which does not accept eight electrons, is very reactive. Information technology readily combines with a molecule containing an atom with a lone pair of electrons. For instance, NH_three reacts with BF_three because the lone pair on nitrogen can be shared with the boron atom:

Hypervalent Molecules

Elements in the second period of the periodic table (northward = 2) can accommodate only eight electrons in their valence crush orbitals because they have only four valence orbitals (one 2s and three 2p orbitals). Elements in the 3rd and higher periods (n ≥ three) have more than four valence orbitals and tin can share more than than 4 pairs of electrons with other atoms considering they have empty d orbitals in the same shell. Molecules formed from these elements are sometimes called hypervalent molecules. Effigy 4 shows the Lewis structures for 2 hypervalent molecules, PCl_v and SF_vi.

Figure 4. In PCl₅, the central atom phosphorus shares five pairs of electrons. In SF_half-dozen, sulfur shares vi pairs of electrons.

In some hypervalent molecules, such as IF₅ and XeF₄, some of the electrons in the outer shell of the central cantlet are lone pairs:

When we write the Lewis structures for these molecules, we discover that we have electrons left over later filling the valence shells of the outer atoms with eight electrons. These additional electrons must be assigned to the fundamental atom.

Example two

Writing Lewis Structures: Octet Rule Violations
Xenon is a noble gas, only it forms a number of stable compounds. Nosotros examined XeF₄ before. What are the Lewis structures of XeF₂ and XeF₆?

Solution
We can draw the Lewis structure of any covalent molecule by following the six steps discussed earlier. In this case, we can condense the final few steps, since not all of them employ.

1. Calculate the number of valence electrons: XeF₂: 8 + (2 × 7) = 22XeF₆: 8 + (6 × 7) = 50
2. Draw a skeleton joining the atoms by single bonds. Xenon will exist the central cantlet because fluorine cannot be a central atom:
   
3. Distribute the remaining electrons.XeF₂: We identify three lone pairs of electrons around each F atom, accounting for 12 electrons and giving each F atom 8 electrons. Thus, six electrons (three alone pairs) remain. These lone pairs must exist placed on the Xe atom. This is acceptable because Xe atoms have empty valence shell d orbitals and tin can accommodate more than eight electrons. The Lewis structure of XeF_two shows two bonding pairs and 3 lone pairs of electrons effectually the Xe atom:
   XeF₆: We place three lone pairs of electrons around each F atom, bookkeeping for 36 electrons. Two electrons remain, and this lone pair is placed on the Xe atom:

Check Your Learning
The halogens form a class of compounds chosen the interhalogens, in which halogen atoms covalently bond to each other. Write the Lewis structures for the interhalogens BrCl₃ and ICl_four ⁻.

Answer:

Key Concepts and Summary

Valence electronic structures can exist visualized past drawing Lewis symbols (for atoms and monatomic ions) and Lewis structures (for molecules and polyatomic ions). Lone pairs, unpaired electrons, and single, double, or triple bonds are used to indicate where the valence electrons are located around each cantlet in a Lewis structure. Almost structures—especially those containing 2nd row elements—obey the octet rule, in which every atom (except H) is surrounded by eight electrons. Exceptions to the octet rule occur for odd-electron molecules (gratuitous radicals), electron-deficient molecules, and hypervalent molecules.

Chemistry End of Chapter Exercises

1. Write the Lewis symbols for each of the following ions:

   (a) As^3–

   (b) I^–

   (c) Beⁱⁱ⁺

   (d) O^ii–

   (e) Ga³⁺

   (f) Li⁺

   (yard) N^three–

2. Many monatomic ions are constitute in seawater, including the ions formed from the post-obit list of elements. Write the Lewis symbols for the monatomic ions formed from the following elements:

   (a) Cl

   (b) Na

   (c) Mg

   (d) Ca

   (e) Grand

   (f) Br

   (g) Sr

   (h) F

3. Write the Lewis symbols of the ions in each of the post-obit ionic compounds and the Lewis symbols of the atom from which they are formed:

   (a) MgS

   (b) Al₂O₃

   (c) GaCl_iii

   (d) K_twoO

   (eastward) Li_threeNorthward

   (f) KF

4. In the Lewis structures listed hither, Grand and Ten correspond various elements in the third period of the periodic table. Write the formula of each compound using the chemical symbols of each element:

   (a)

   

   (b)

   

   (c)

   

   (d)

   

5. Write the Lewis structure for the diatomic molecule P₂, an unstable grade of phosphorus found in loftier-temperature phosphorus vapor.
6. Write Lewis structures for the post-obit:

   (a) H₂

   (b) HBr

   (c) PCl₃

   (d) SF_ii

   (east) H₂CCH₂

   (f) HNNH

   (one thousand) H_iiCNH

   (h) NO^–

   (i) N_ii

   (j) CO

   (thousand) CN^–

7. Write Lewis structures for the following:

   (a) O₂

   (b) H_iiCO

   (c) AsF_iii

   (d) ClNO

   (e) SiCl_iv

   (f) H₃O⁺

   (g) NH₄ ⁺

   (h) BF₄ ⁻

   (i) HCCH

   (j) ClCN

   (k) C₂ ²⁺

8. Write Lewis structures for the following:

   (a) ClF_iii

   (b) PCl₅

   (c) BF_three

   (d) PF₆ ⁻

9. Write Lewis structures for the following:

   (a) SeF₆

   (b) XeF₄

   (c) SeCl₃ ⁺

   (d) Cl₂BBCl_two (contains a B–B bail)

10. Write Lewis structures for:

    (a) PO_four ^three−

    (b) ICl₄ ⁻

    (c) Then₃ ²⁻

    (d) HONO

11. Correct the post-obit argument: "The bonds in solid PbCl_ii are ionic; the bond in a HCl molecule is covalent. Thus, all of the valence electrons in PbCl_two are located on the Cl^– ions, and all of the valence electrons in a HCl molecule are shared between the H and Cl atoms."
12. Write Lewis structures for the following molecules or ions:

    (a) SbH_iii

    (b) XeF₂

    (c) Se₈ (a cyclic molecule with a ring of eight Se atoms)

13. Methanol, H₃COH, is used as the fuel in some race cars. Ethanol, C₂H₅OH, is used extensively every bit motor fuel in Brazil. Both methanol and ethanol produce CO₂ and H₂O when they burn. Write the chemical equations for these combustion reactions using Lewis structures instead of chemical formulas.
14. Many planets in our solar arrangement contain organic chemicals including methane (CH₄) and traces of ethylene (C₂H₄), ethane (C₂H₆), propyne (H₃CCCH), and diacetylene (HCCCCH). Write the Lewis structures for each of these molecules.
15. Carbon tetrachloride was formerly used in fire extinguishers for electrical fires. It is no longer used for this purpose because of the formation of the toxic gas phosgene, Cl_iiCO. Write the Lewis structures for carbon tetrachloride and phosgene.
16. Identify the atoms that correspond to each of the following electron configurations. So, write the Lewis symbol for the common ion formed from each cantlet:

    (a) 1south ²iisouthward ²2p ⁵

    (b) anes ^two2south ²2p ⁶3s ²

    (c) 1s ²2south ^twotwop ^six3s ²iiip ^viivs ⁱⁱ3d ¹⁰

    (d) 1s ²2s ²2p ^{half dozen}threes ²3p ^{half dozen}ivs ²3d ¹⁰ivp ^iv

    (e) 1southward ²2s ⁱⁱiip ^{half dozen}3south ²3p ⁶4s ^two3d ¹⁰4p ¹

17. The arrangement of atoms in several biologically important molecules is given here. Consummate the Lewis structures of these molecules by adding multiple bonds and lone pairs. Practice non add any more atoms.

    (a) the amino acid serine:

    

    (b) urea:

    

    (c) pyruvic acrid:

    

    (d) uracil:

    

    (e) carbonic acrid:

    

18. A compound with a molar mass of about 28 g/mol contains 85.seven% carbon and 14.three% hydrogen by mass. Write the Lewis structure for a molecule of the compound.
19. A chemical compound with a molar mass of about 42 g/mol contains 85.seven% carbon and 14.iii% hydrogen by mass. Write the Lewis structure for a molecule of the compound.
20. Two arrangements of atoms are possible for a chemical compound with a tooth mass of about 45 thousand/mol that contains 52.two% C, 13.i% H, and 34.7% O by mass. Write the Lewis structures for the ii molecules.
21. How are single, double, and triple bonds similar? How practice they differ?

Glossary

double bond

covalent bail in which ii pairs of electrons are shared between two atoms

free radical

molecule that contains an odd number of electrons

hypervalent molecule

molecule containing at least one principal group chemical element that has more eight electrons in its valence shell

Lewis structure

diagram showing lone pairs and bonding pairs of electrons in a molecule or an ion

Lewis symbol

symbol for an chemical element or monatomic ion that uses a dot to stand for each valence electron in the chemical element or ion

lone pair

2 (a pair of) valence electrons that are not used to form a covalent bond

octet rule

guideline that states main group atoms will course structures in which eight valence electrons interact with each nucleus, counting bonding electrons as interacting with both atoms connected by the bond

unmarried bail

bond in which a unmarried pair of electrons is shared between 2 atoms

triple bail

bail in which three pairs of electrons are shared between two atoms

Solutions

Answers to Chemistry End of Chapter Exercises

ane. (a) viii electrons:
;

(b) eight electrons:

;

(c) no electrons

Exist²⁺;

(d) 8 electrons:

;

(e) no electrons

Ga³⁺;

(f) no electrons

Li⁺;

(g) eight electrons:

3. (a)

;

(b)

;

(c)

;

(d)

>;

(e)

;

(f)

5.

vii. (a)

In this instance, the Lewis construction is inadequate to draw the fact that experimental studies have shown two unpaired electrons in each oxygen molecule.

(b)

;

(c)

;

(d)

;

(e)

;

(f)

;

(thou)

;

(h)

;

(i)

;

(j)

;

(k)

9. (a) SeF_half-dozen:
;

(b) XeF₄:

;

(c) SeCl₃ ⁺:

;

(d) Cl₂BBCl₂:

11. 2 valence electrons per Lead atom are transferred to Cl atoms; the resulting Pb²⁺ ion has a 6s ^two valence trounce configuration. Two of the valence electrons in the HCl molecule are shared, and the other half dozen are located on the Cl atom as lone pairs of electrons.

13.

15.

17. (a)
;

(b)

;

(c)

;

(d)

;

(e)

xix.

21. Each bail includes a sharing of electrons between atoms. Two electrons are shared in a single bond; four electrons are shared in a double bond; and six electrons are shared in a triple bond.

Source: https://opentextbc.ca/chemistry/chapter/7-3-lewis-symbols-and-structures/

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