Write a Generic Lewis Structure for the Elements in the Oxygen Family (Group 6a).

Chapter 7. Chemic Bonding and Molecular Geometry

seven.3 Lewis Symbols and Structures

Learning Objectives

By the cease of this section, you volition be able to:

• Write Lewis symbols for neutral atoms and ions
• Draw Lewis structures depicting the bonding in simple molecules

Thus far in this chapter, we 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 trounce electrons betwixt atoms. In this section, we volition explore the typical method for depicting valence shell electrons and chemical bonds, namely Lewis symbols and Lewis structures.

Lewis Symbols

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

Figure ane shows the Lewis symbols for the elements of the tertiary flow of the periodic table.

Effigy 1. Lewis symbols illustrating the number of valence electrons for each element in the third period of the periodic tabular array.

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

As well, they can be used to show the formation of anions from atoms, equally shown hither for chlorine and sulfur:

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

Effigy 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 not change.

Lewis Structures

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

The Lewis structure indicates that each Cl cantlet has three pairs of electrons that are not used in bonding (chosen lone pairs) and ane 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 called a single bond. Each Cl atom interacts with viii valence electrons: the six in the lone pairs and the two in the single bond.

The Octet Dominion

The other halogen molecules (F₂, Br₂, I₂, and At_two) course bonds like those in the chlorine molecule: 1 single bond between atoms and three lonely pairs of electrons per atom. This allows each halogen atom to take a noble gas electron configuration. The tendency of main group atoms to grade enough bonds to obtain eight valence electrons is known as the octet rule.

The number of bonds that an atom tin form can frequently be predicted from the number of electrons needed to reach an octet (viii valence electrons); this is especially true of the nonmetals of the second flow of the periodic table (C, N, O, and F). For example, each atom of a group fourteen element has four electrons in its outermost shell and therefore requires four more than electrons to reach an octet. These four electrons tin be gained past forming four covalent bonds, equally illustrated here for carbon in CCl_four (carbon tetrachloride) and silicon in SiH_iv (silane). Because hydrogen only needs two electrons to fill up its valence beat out, it is an exception to the octet rule. The transition elements and inner transition elements also do not follow the octet rule:

Grouping 15 elements such every bit nitrogen have five valence electrons in the diminutive Lewis symbol: ane lone pair and three unpaired electrons. To obtain an octet, these atoms form three covalent bonds, as in NH_three (ammonia). Oxygen and other atoms in grouping xvi obtain an octet by forming two covalent bonds:

Double and Triple Bonds

As previously mentioned, when a pair of atoms shares i pair of electrons, we call this a single bond. However, a pair of atoms may need to share more than than 1 pair of electrons in order to accomplish the requisite octet. A double bond forms when ii pairs of electrons are shared between a pair of atoms, as between the carbon and oxygen atoms in CH_twoO (formaldehyde) and betwixt the two carbon atoms in C₂H_iv (ethylene):

A triple bail 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 unproblematic molecules and molecular ions, we tin can write the Lewis structures past merely pairing upwards the unpaired electrons on the elective atoms. Meet these examples:

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

1. Determine the total number of valence (outer shell) electrons. For cations, subtract i electron for each positive charge. For anions, add together one electron for each negative charge.
2. Draw a skeleton structure of the molecule or ion, arranging the atoms effectually a cardinal atom. (Mostly, the least electronegative chemical element should exist placed in the center.) Connect each atom to the central atom with a single bond (one electron pair).
3. Distribute the remaining electrons as lone pairs on the terminal atoms (except hydrogen), completing an octet effectually each cantlet.
4. Place all remaining electrons on the central atom.
5. Rearrange the electrons of the outer atoms to make multiple bonds with the central atom in order to obtain octets wherever possible.

Let us determine the Lewis structures of SiH_four, CHO_two−, NO⁺, and OF₂ as examples in following this procedure:

1. Determine the full number of valence (outer beat out) electrons in the molecule or ion.
   • For a molecule, we add together the number of valence electrons on each cantlet in the molecule:

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

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

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

   • For a positive ion, such as NO⁺, nosotros add the number of valence electrons on the atoms in the ion and and so subtract the number of positive charges on the ion (ane electron is lost for each single positive accuse) from the total number of valence electrons:

     [latex]\begin{array}{r r 50} \text{NO}^{+} & & \\[1em] & \text{Due north: 5 valence electrons/cantlet} \times one \;\text{atom} & = 5 \\[1em] & \text{O: vi valence electrons/atom} \times one \;\text{atom} & = 6 \\[1em] \dominion[-0.5ex]{21em}{0.1ex}\hspace{-21em} + & -1 \;\text{electron (positive accuse)} & = -1 \\[1em] & & = 10 \;\text{valence electrons} \end{assortment}[/latex]

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

     [latex]\begin{array}{r r 50} \text{OF}_{2} & & \\[1em] & \text{O: 6 valence electrons/atom} \times 1 \;\text{atom} & = half dozen \\[1em] \rule[-0.5ex]{21em}{0.1ex}\hspace{-21em} + & \text{F: 7 valence electrons/atom} \times 2 \;\text{atoms} & = xiv \\[1em] & & = 20 \;\text{valence electrons} \finish{array}[/latex]

2. Draw a skeleton construction of the molecule or ion, arranging the atoms effectually a central cantlet and connecting each cantlet to the cardinal atom with a unmarried (one electron pair) bail. (Notation that nosotros denote ions with brackets around the structure, indicating the charge outside the brackets:)When several arrangements of atoms are possible, as for CHO_ii ⁻, we must use experimental prove to choose the right i. In general, the less electronegative elements are more than probable 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 fundamental atom. As the near electronegative element, fluorine besides cannot exist a central cantlet.
3. Distribute the remaining electrons as lone pairs on the terminal atoms (except hydrogen) to complete their valence shells with an octet of electrons.
   • In that location are no remaining electrons on SiH₄, then information technology is unchanged:
4. Place all remaining electrons on the primal atom.
   • For SiH₄, CHO_ii ⁻, and NO⁺, in that location are no remaining electrons; we already placed all of the electrons determined in Stride 1.
   • For OF_ii, we had 16 electrons remaining in Step iii, and we placed 12, leaving 4 to exist placed on the key atom:
5. Rearrange the electrons of the outer atoms to brand multiple bonds with the fundamental atom in order to obtain octets wherever possible.

Example i

Writing Lewis Structures
NASA's Cassini-Huygens mission detected a large deject of toxic hydrogen cyanide (HCN) on Titan, 1 of Saturn's moons. Titan as well contains ethane (H₃CCH₃), acetylene (HCCH), and ammonia (NH₃). What are the Lewis structures of these molecules?

Solution

1. Calculate the number of valence electrons.HCN: (i × 1) + (4 × i) + (5 × ane) = 10H_iiiCCH₃: (ane × 3) + (2 × 4) + (i × iii) = 14HCCH: (1 × 1) + (2 × four) + (ane × 1) = 10NH₃: (5 × i) + (3 × 1) = viii
2. Draw a skeleton and connect the atoms with single bonds. Call up that H is never a fundamental atom:
3. Where needed, distribute electrons to the concluding atoms: HCN: 6 electrons placed on NH₃CCH₃: no electrons remainHCCH: no concluding atoms capable of accepting electrons

   NH_iii: no terminal atoms capable of accepting electrons

4. Where needed, identify remaining electrons on the central atom: HCN: no electrons remainH_threeCCH₃: no electrons remainHCCH: four electrons placed on carbon

   NH_three: two electrons placed on nitrogen

5. Where needed, rearrange electrons to course multiple bonds in social club to obtain an octet on each atom:HCN: form 2 more C–N bondsH₃CCH₃: all atoms accept the correct number of electronsHCCH: form a triple bond between the two carbon atomsNH₃: all atoms have the correct number of electrons

   

Check Your Learning
Both carbon monoxide, CO, and carbon dioxide, CO₂, are products of the combustion of fossil fuels. Both of these gases as well cause problems: CO is toxic and CO_two has been implicated in global climate alter. What are the Lewis structures of these ii molecules?

Answer:

Fullerene Chemistry

Carbon soot has been known to homo since prehistoric times, but information technology was not until adequately recently that the molecular structure of the main component of soot was discovered. In 1996, the Nobel Prize in Chemical science was awarded to Richard Smalley (Figure 3), Robert Gyre, and Harold Kroto for their work in discovering a new form of carbon, the C_threescore buckminsterfullerene molecule (Figure 1 in Chapter 7 Introduction). An entire course of compounds, including spheres and tubes of various shapes, were discovered based on C_sixty. This type of molecule, called a fullerene, shows promise in a diversity of applications. Considering of their size and shape, fullerenes tin encapsulate other molecules, so they have shown potential in various applications from hydrogen storage to targeted drug delivery systems. They also possess unique electronic and optical properties that have been put to good use in solar powered devices and chemical sensors.

Figure 3. Richard Smalley (1943–2005), a professor of physics, chemistry, and astronomy at Rice University, was one of the leading advocates for fullerene chemistry. Upon his death in 2005, the US Senate honored him as the "Father of Nanotechnology." (credit: U.s.a. Section of Energy)

Exceptions to the Octet Rule

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

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

Odd-electron Molecules

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

To draw the Lewis construction for an odd-electron molecule similar NO, we follow the aforementioned five steps we would for other molecules, just with a few minor changes:

1. Determine the full number of valence (outer crush) electrons. The sum of the valence electrons is 5 (from N) + vi (from O) = eleven. The odd number immediately tells us that nosotros have a free radical, and then we know that not every atom tin can take eight electrons in its valence shell.
2. Draw a skeleton structure of the molecule. Nosotros tin easily draw a skeleton with an North–O single bail:N–O
3. Distribute the remaining electrons as solitary pairs on the concluding atoms. In this case, there is no central atom, so nosotros distribute the electrons effectually both atoms. We give viii electrons to the more electronegative atom in these situations; thus oxygen has the filled valence beat:
   
4. Place all remaining electrons on the fundamental atom. Since there are no remaining electrons, this step does non apply.
5. Rearrange the electrons to make multiple bonds with the central atom in order to obtain octets wherever possible. Nosotros know that an odd-electron molecule cannot accept an octet for every cantlet, but nosotros desire to get each atom every bit close to an octet every bit possible. In this case, nitrogen has only five electrons around information technology. To move closer to an octet for nitrogen, we take ane of the lone pairs from oxygen and utilise it to grade a NO double bond. (We cannot take another lonely pair of electrons on oxygen and form a triple bond because nitrogen would then have 9 electrons:)
   

Electron-deficient Molecules

Nosotros will also meet a few molecules that contain central atoms that do not have a filled valence shell. By and large, these are molecules with central atoms from groups ii and 12, outer atoms that are hydrogen, or other atoms that practice not form multiple bonds. For case, in the Lewis structures of glucinium dihydride, BeH₂, and boron trifluoride, BF₃, the beryllium and boron atoms each have simply 4 and six 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, but experimental evidence indicates the bail lengths are closer to that expected for B–F unmarried bonds. This suggests the best Lewis structure has three B–F single bonds and an electron deficient boron. The reactivity of the compound is also consistent with an electron deficient boron. Even so, the B–F bonds are slightly shorter than what is really expected for B–F unmarried bonds, indicating that some double bond graphic symbol is found in the bodily molecule.

An atom like the boron atom in BF₃, which does non have eight electrons, is very reactive. It readily combines with a molecule containing an atom with a lonely pair of electrons. For instance, NH₃ reacts with BF_iii because the lone pair on nitrogen tin be shared with the boron atom:

Hypervalent Molecules

Elements in the second period of the periodic table (n = ii) tin arrange only eight electrons in their valence shell orbitals because they have only four valence orbitals (i 2s and three 2p orbitals). Elements in the third and higher periods (n ≥ 3) have more than four valence orbitals and can share more four pairs of electrons with other atoms because they have empty d orbitals in the aforementioned shell. Molecules formed from these elements are sometimes called hypervalent molecules. Figure iv shows the Lewis structures for two hypervalent molecules, PCl₅ and SF_6.

Figure 4. In PCl₅, the fundamental atom phosphorus shares 5 pairs of electrons. In SF_half-dozen, sulfur shares six pairs of electrons.

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

When we write the Lewis structures for these molecules, we find that nosotros take electrons left over after filling the valence shells of the outer atoms with eight electrons. These additional electrons must exist assigned to the fundamental atom.

Example 2

Writing Lewis Structures: Octet Rule Violations
Xenon is a noble gas, just it forms a number of stable compounds. We examined XeF₄ earlier. What are the Lewis structures of XeF_ii and XeF₆?

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

1. Calculate the number of valence electrons: XeF₂: 8 + (two × 7) = 22XeF₆: 8 + (6 × 7) = fifty
2. Describe a skeleton joining the atoms past single bonds. Xenon will be the central atom because fluorine cannot be a cardinal atom:
   
3. Distribute the remaining electrons.XeF₂: We place three lone pairs of electrons effectually each F atom, accounting for 12 electrons and giving each F cantlet 8 electrons. Thus, half dozen electrons (three solitary pairs) remain. These solitary pairs must be placed on the Xe cantlet. This is adequate because Xe atoms have empty valence beat d orbitals and can accommodate more than eight electrons. The Lewis structure of XeF₂ shows two bonding pairs and iii lone pairs of electrons effectually the Xe cantlet:
   XeF_{half dozen}: We place three alone pairs of electrons around each F cantlet, bookkeeping for 36 electrons. Two electrons remain, and this lone pair is placed on the Xe atom:

Cheque Your Learning
The halogens form a class of compounds called the interhalogens, in which halogen atoms covalently bond to each other. Write the Lewis structures for the interhalogens BrCl₃ and ICl₄ ⁻.

Respond:

Key Concepts and Summary

Valence electronic structures tin can exist visualized by 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 atom in a Lewis construction. Most structures—peculiarly those containing second 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 Cease of Chapter Exercises

1. Write the Lewis symbols for each of the post-obit ions:

   (a) As^three–

   (b) I^–

   (c) Be²⁺

   (d) O^ii–

   (due east) Ga³⁺

   (f) Li⁺

   (g) N^3–

2. Many monatomic ions are found 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

   (eastward) K

   (f) Br

   (one thousand) Sr

   (h) F

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

   (a) MgS

   (b) Al₂O₃

   (c) GaCl₃

   (d) 1000₂O

   (e) Li_iiiN

   (f) KF

4. In the Lewis structures listed hither, M and X represent various elements in the third menstruum of the periodic tabular array. 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 form of phosphorus found in high-temperature phosphorus vapor.
6. Write Lewis structures for the following:

   (a) H₂

   (b) HBr

   (c) PCl₃

   (d) SF_two

   (e) H₂CCH₂

   (f) HNNH

   (chiliad) H₂CNH

   (h) NO^–

   (i) N_two

   (j) CO

   (m) CN^–

7. Write Lewis structures for the post-obit:

   (a) O_two

   (b) H_iiCO

   (c) AsF_iii

   (d) ClNO

   (due east) SiCl_four

   (f) H₃O⁺

   (g) NH_four ⁺

   (h) BF_four ⁻

   (i) HCCH

   (j) ClCN

   (k) C₂ ^two+

8. Write Lewis structures for the following:

   (a) ClF₃

   (b) PCl₅

   (c) BF_iii

   (d) PF₆ ⁻

9. Write Lewis structures for the following:

   (a) SeF_{half dozen}

   (b) XeF₄

   (c) SeCl_three ⁺

   (d) Cl₂BBCl₂ (contains a B–B bond)

10. Write Lewis structures for:

    (a) PO₄ ³⁻

    (b) ICl_four ⁻

    (c) SO₃ ²⁻

    (d) HONO

11. Correct the following statement: "The bonds in solid PbCl_two are ionic; the bond in a HCl molecule is covalent. Thus, all of the valence electrons in PbCl₂ 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) SbH₃

    (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 as motor fuel in Brazil. Both methanol and ethanol produce CO₂ and H₂O when they fire. Write the chemical equations for these combustion reactions using Lewis structures instead of chemical formulas.
14. Many planets in our solar system comprise organic chemicals including methyl hydride (CH₄) and traces of ethylene (C₂H₄), ethane (C₂H₆), propyne (H_iiiCCCH), 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 germination of the toxic gas phosgene, Cl₂CO. Write the Lewis structures for carbon tetrachloride and phosgene.
16. Identify the atoms that correspond to each of the post-obit electron configurations. Then, write the Lewis symbol for the common ion formed from each cantlet:

    (a) 1s ²iidue south ²twop ⁵

    (b) 1s ⁱⁱ2s ²2p ⁶3s ²

    (c) is ²iisouthward ²2p ⁶3s ²3p ⁶4south ²3d ¹⁰

    (d) 1s ²2southward ²2p ⁶3s ²threep ^{half dozen}ivs ²3d ¹⁰ivp ⁴

    (due east) 1south ⁱⁱ2due south ⁱⁱ2p ⁶3south ⁱⁱiiip ⁶4south ^twothreed ¹⁰ivp ¹

17. The system of atoms in several biologically of import molecules is given hither. Consummate the Lewis structures of these molecules by adding multiple bonds and lonely pairs. Do not add any more atoms.

    (a) the amino acid serine:

    

    (b) urea:

    

    (c) pyruvic acrid:

    

    (d) uracil:

    

    (eastward) carbonic acrid:

    

18. A compound with a molar mass of near 28 g/mol contains 85.7% carbon and fourteen.three% hydrogen by mass. Write the Lewis structure for a molecule of the chemical compound.
19. A compound with a molar mass of about 42 g/mol contains 85.7% carbon and 14.3% hydrogen past mass. Write the Lewis structure for a molecule of the compound.
20. Two arrangements of atoms are possible for a compound with a tooth mass of virtually 45 thou/mol that contains 52.2% C, xiii.1% H, and 34.7% O by mass. Write the Lewis structures for the two molecules.
21. How are unmarried, double, and triple bonds similar? How exercise they differ?

Glossary

double bond

covalent bond in which 2 pairs of electrons are shared between ii atoms

gratis radical

molecule that contains an odd number of electrons

hypervalent molecule

molecule containing at least one main grouping element that has more than eight electrons in its valence shell

Lewis structure

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

Lewis symbol

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

lone pair

two (a pair of) valence electrons that are not used to form a covalent bail

octet rule

guideline that states principal grouping atoms will form structures in which eight valence electrons interact with each nucleus, counting bonding electrons as interacting with both atoms continued by the bond

unmarried bond

bail in which a single pair of electrons is shared between two atoms

triple bond

bond in which three pairs of electrons are shared between ii atoms

Solutions

Answers to Chemistry End of Chapter Exercises

1. (a) viii electrons:
;

(b) eight electrons:

;

(c) no electrons

Be²⁺;

(d) eight electrons:

;

(due east) no electrons

Ga³⁺;

(f) no electrons

Li⁺;

(g) viii electrons:

3. (a)

;

(b)

;

(c)

;

(d)

>;

(due east)

;

(f)

5.

7. (a)

In this case, the Lewis structure is inadequate to depict the fact that experimental studies have shown two unpaired electrons in each oxygen molecule.

(b)

;

(c)

;

(d)

;

(east)

;

(f)

;

(g)

;

(h)

;

(i)

;

(j)

;

(thousand)

9. (a) SeF₆:
;

(b) XeF₄:

;

(c) SeCl₃ ⁺:

;

(d) Cl_twoBBCl_two:

11. 2 valence electrons per Pb atom are transferred to Cl atoms; the resulting Pb²⁺ ion has a 6southward ² valence crush configuration. Two of the valence electrons in the HCl molecule are shared, and the other six are located on the Cl atom equally solitary pairs of electrons.

thirteen.

15.

17. (a)
;

(b)

;

(c)

;

(d)

;

(e)

19.

21. Each bond includes a sharing of electrons between atoms. 2 electrons are shared in a single bond; four electrons are shared in a double bond; and six electrons are shared in a triple bond.

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