How To Determine The Base Element In A Covalent Bond When Drawing A Lewis Dot

Chapter seven. Chemic Bonding and Molecular Geometry

7.iii Lewis Symbols and Structures

Learning Objectives

Past the terminate of this department, y'all will 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 take discussed the various types of bonds that grade between atoms and/or ions. In all cases, these bonds involve the sharing or transfer of valence shell electrons between atoms. In this department, we will explore the typical method for depicting valence crush electrons and chemical bonds, namely Lewis symbols and Lewis structures.

Lewis Symbols

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

Figure 1 shows the Lewis symbols for the elements of the third menstruation of the periodic table.

Figure i. Lewis symbols illustrating the number of valence electrons for each element in the third catamenia of the periodic tabular array.

Lewis symbols tin can besides be used to illustrate the formation of cations from atoms, every bit shown here for sodium and calcium:

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

Effigy 2 demonstrates the employ of Lewis symbols to show 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 by atoms gaining electrons. The full number of electrons does not change.

Lewis Structures

We as well use Lewis symbols to indicate the germination 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 class a chlorine molecule, they share i pair of electrons:

The Lewis structure indicates that each Cl cantlet has iii pairs of electrons that are not used in bonding (called lone pairs) and ane shared pair of electrons (written betwixt the atoms). A dash (or line) is sometimes used to point a shared pair of electrons:

A single shared pair of electrons is called a single bail. Each Cl atom interacts with eight valence electrons: the six in the lone pairs and the two in the single bail.

The Octet Dominion

The other halogen molecules (F_ii, Br_two, I₂, and At₂) form bonds similar those in the chlorine molecule: 1 single bail between atoms and three lone pairs of electrons per atom. This allows each halogen cantlet to have a noble gas electron configuration. The trend of master grouping atoms to grade plenty bonds to obtain eight valence electrons is known as the octet rule.

The number of bonds that an atom can class can often exist predicted from the number of electrons needed to reach an octet (viii valence electrons); this is especially true of the nonmetals of the 2d menstruum of the periodic table (C, N, O, and F). For case, each atom of a group 14 element has four electrons in its outermost shell and therefore requires four more electrons to reach an octet. These four electrons tin can exist gained by forming four covalent bonds, as illustrated here for carbon in CCl₄ (carbon tetrachloride) and silicon in SiH_four (silane). Considering hydrogen only needs two electrons to fill its valence shell, it is an exception to the octet rule. The transition elements and inner transition elements also practice not follow the octet rule:

Grouping 15 elements such as nitrogen have five valence electrons in the atomic Lewis symbol: 1 lone pair and three unpaired electrons. To obtain an octet, these atoms class three covalent bonds, as in NH_three (ammonia). Oxygen and other atoms in group 16 obtain an octet by forming two covalent bonds:

Double and Triple Bonds

As previously mentioned, when a pair of atoms shares one pair of electrons, we phone call this a unmarried bond. However, a pair of atoms may need to share more than ane pair of electrons in society to attain 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 CH₂O (formaldehyde) and between the 2 carbon atoms in C₂H₄ (ethylene):

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

Writing Lewis Structures with the Octet Rule

For very unproblematic molecules and molecular ions, we can write the Lewis structures by simply pairing upwards the unpaired electrons on the elective atoms. Encounter these examples:

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

1. Make up one's mind the full number of valence (outer crush) electrons. For cations, decrease one electron for each positive charge. For anions, add one electron for each negative charge.
2. Draw a skeleton structure of the molecule or ion, arranging the atoms around a central atom. (More often than not, the least electronegative element should be placed in the center.) Connect each atom to the cardinal atom with a unmarried bail (ane electron pair).
3. Distribute the remaining electrons as lonely pairs on the terminal atoms (except hydrogen), completing an octet effectually each atom.
4. Place all remaining electrons on the central cantlet.
5. Rearrange the electrons of the outer atoms to brand multiple bonds with the primal atom in lodge to obtain octets wherever possible.

Let us determine the Lewis structures of SiH₄, CHO₂−, NO⁺, and OF₂ equally examples in following this procedure:

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

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

   • For a negative ion, such as CHO_two ⁻, 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 single negative accuse):

     [latex]\begin{assortment}{r r fifty} {\text{CHO}_2}^{-} & & \\[1em] & \text{C: 4 valence electrons/atom} \times i \;\text{atom} & = 4 \\[1em] & \text{H: 1 valence electron/cantlet} \times 1 \;\text{atom} & = i \\[1em] & \text{O: half dozen valence electrons/atom} \times 2 \;\text{atoms} & = 12 \\[1em] \rule[-0.5ex]{21.5em}{0.1ex}\hspace{-21.5em} + & 1\;\text{additional electron} & = 1 \\[1em] & & = 18 \;\text{valence electrons} \end{assortment}[/latex]

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

     [latex]\begin{array}{r r l} \text{NO}^{+} & & \\[1em] & \text{N: 5 valence electrons/cantlet} \times ane \;\text{atom} & = v \\[1em] & \text{O: six valence electrons/atom} \times i \;\text{cantlet} & = 6 \\[1em] \rule[-0.5ex]{21em}{0.1ex}\hspace{-21em} + & -1 \;\text{electron (positive charge)} & = -1 \\[1em] & & = 10 \;\text{valence electrons} \end{array}[/latex]

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

     [latex]\begin{array}{r r fifty} \text{OF}_{2} & & \\[1em] & \text{O: 6 valence electrons/atom} \times 1 \;\text{cantlet} & = half dozen \\[1em] \dominion[-0.5ex]{21em}{0.1ex}\hspace{-21em} + & \text{F: seven valence electrons/cantlet} \times 2 \;\text{atoms} & = 14 \\[1em] & & = 20 \;\text{valence electrons} \stop{array}[/latex]

2. Describe a skeleton structure of the molecule or ion, arranging the atoms around a central atom and connecting each atom to the central atom with a single (one electron pair) bond. (Note that nosotros denote ions with brackets around the structure, indicating the charge outside the brackets:)When several arrangements of atoms are possible, as for CHO₂ ⁻, we must use experimental bear witness to choose the correct one. 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 information technology. Other examples include P in POCl_iii, S in And so₂, and Cl in ClO₄ ⁻. An exception is that hydrogen is nearly never a primal atom. Equally the about electronegative element, fluorine too cannot be a central atom.
3. Distribute the remaining electrons as lone pairs on the terminal atoms (except hydrogen) to complete their valence shells with an octet of electrons.
   • There are no remaining electrons on SiH₄, so it is unchanged:
4. Identify all remaining electrons on the central atom.
   • For SiH₄, CHO_ii ⁻, and NO⁺, in that location are no remaining electrons; we already placed all of the electrons determined in Footstep one.
   • For OF_two, we had 16 electrons remaining in Step 3, and nosotros placed 12, leaving iv to be placed on the primal cantlet:
5. Rearrange the electrons of the outer atoms to make multiple bonds with the fundamental cantlet in guild to obtain octets wherever possible.

Example i

Writing Lewis Structures
NASA'south Cassini-Huygens mission detected a large cloud of toxic hydrogen cyanide (HCN) on Titan, i of Saturn's moons. Titan also 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: (1 × 1) + (4 × 1) + (v × one) = 10H₃CCH₃: (one × iii) + (2 × 4) + (1 × 3) = 14HCCH: (1 × ane) + (2 × four) + (1 × 1) = 10NH_iii: (5 × 1) + (3 × 1) = 8
2. Describe a skeleton and connect the atoms with single bonds. Remember that H is never a central atom:
3. Where needed, distribute electrons to the last atoms: HCN: six electrons placed on NH₃CCH_three: no electrons remainHCCH: no terminal atoms capable of accepting electrons

   NH₃: no final atoms capable of accepting electrons

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

   NH₃: 2 electrons placed on nitrogen

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

   

Bank check Your Learning
Both carbon monoxide, CO, and carbon dioxide, CO₂, are products of the combustion of fossil fuels. Both of these gases too cause problems: CO is toxic and CO_ii has been implicated in global climate change. What are the Lewis structures of these two molecules?

Respond:

Fullerene Chemical science

Carbon soot has been known to human being since prehistoric times, merely it was not 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 three), Robert Curl, and Harold Kroto for their work in discovering a new class of carbon, the C_sixty buckminsterfullerene molecule (Figure 1 in Chapter 7 Introduction). An entire class 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 variety of applications. Because 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 take been put to good use in solar powered devices and chemic sensors.

Effigy 3. Richard Smalley (1943–2005), a professor of physics, chemistry, and astronomy at Rice Academy, was one of the leading advocates for fullerene chemistry. Upon his expiry in 2005, the U.s. Senate honored him as the "Father of Nanotechnology." (credit: The states Department of Energy)

Exceptions to the Octet Dominion

Many covalent molecules take central atoms that exercise not accept viii electrons in their Lewis structures. These molecules autumn into three categories:

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

Odd-electron Molecules

Nosotros call molecules that contain 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 structure for an odd-electron molecule similar NO, we follow the same five steps nosotros would for other molecules, but with a few small-scale changes:

1. Make up one's mind the total number of valence (outer shell) electrons. The sum of the valence electrons is five (from N) + 6 (from O) = 11. The odd number immediately tells us that we have a free radical, so nosotros know that not every atom tin take 8 electrons in its valence crush.
2. Draw a skeleton structure of the molecule. Nosotros can hands draw a skeleton with an Northward–O unmarried bond:N–O
3. Distribute the remaining electrons as lone pairs on the terminal atoms. In this example, there is no primal atom, so we distribute the electrons around both atoms. We give eight electrons to the more than electronegative atom in these situations; thus oxygen has the filled valence shell:
   
4. Place all remaining electrons on the central atom. Since in that location are no remaining electrons, this pace does not apply.
5. Rearrange the electrons to brand multiple bonds with the primal atom in club to obtain octets wherever possible. We know that an odd-electron molecule cannot accept an octet for every cantlet, just we want to go each atom as close to an octet as possible. In this case, nitrogen has merely five electrons around it. To move closer to an octet for nitrogen, we have one of the lone pairs from oxygen and use information technology to class a NO double bond. (Nosotros cannot have another lone pair of electrons on oxygen and form a triple bail because nitrogen would and so accept nine electrons:)
   

Electron-deficient Molecules

We will also run into a few molecules that contain fundamental atoms that practise non have a filled valence shell. By and large, 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 the Lewis structures of beryllium dihydride, BeH₂, and boron trifluoride, BF₃, the beryllium and boron atoms each have only four and half-dozen electrons, respectively. It is possible to depict a structure with a double bond betwixt a boron cantlet and a fluorine atom in BF_iii, satisfying the octet rule, but experimental evidence indicates the bail lengths are closer to that expected for B–F single bonds. This suggests the best Lewis structure has three B–F unmarried bonds and an electron scarce boron. The reactivity of the compound is also consequent with an electron deficient boron. All the same, the B–F bonds are slightly shorter than what is actually expected for B–F unmarried bonds, indicating that some double bond graphic symbol is constitute in the bodily molecule.

An atom like the boron atom in BF_iii, which does not take viii electrons, is very reactive. It readily combines with a molecule containing an atom with a lone pair of electrons. For case, NH_iii reacts with BF₃ because the lone pair on nitrogen can be shared with the boron atom:

Hypervalent Molecules

Elements in the 2nd period of the periodic tabular array (n = 2) can accommodate only eight electrons in their valence shell orbitals because they have just four valence orbitals (one 2s and iii 2p orbitals). Elements in the 3rd and higher periods (n ≥ 3) take more than than four valence orbitals and tin share more than four pairs of electrons with other atoms considering they accept empty d orbitals in the same shell. Molecules formed from these elements are sometimes called hypervalent molecules. Figure 4 shows the Lewis structures for ii hypervalent molecules, PCl₅ and SF_six.

Figure four. In PCl₅, the central cantlet phosphorus shares five pairs of electrons. In SF₆, sulfur shares six pairs of electrons.

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

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

Example 2

Writing Lewis Structures: Octet Rule Violations
Xenon is a element of group 0, but information technology forms a number of stable compounds. We examined XeF_four before. What are the Lewis structures of XeF_ii and XeF₆?

Solution
We can draw the Lewis construction of any covalent molecule past following the six steps discussed earlier. In this case, we can condense the final few steps, since non all of them use.

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

Bank check Your Learning
The halogens form a class of compounds called the interhalogens, in which element of group vii atoms covalently bail to each other. Write the Lewis structures for the interhalogens BrCl₃ and ICl_iv ⁻.

Answer:

Key Concepts and Summary

Valence electronic structures tin 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 atom in a Lewis structure. Virtually structures—peculiarly those containing second row elements—obey the octet rule, in which every cantlet (except H) is surrounded by eight electrons. Exceptions to the octet dominion occur for odd-electron molecules (gratis radicals), electron-deficient molecules, and hypervalent molecules.

Chemistry End of Chapter Exercises

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

   (a) Every bit^iii–

   (b) I^–

   (c) Be²⁺

   (d) O^2–

   (e) Ga^three+

   (f) Li⁺

   (grand) Due north^3–

2. Many monatomic ions are found in seawater, including the ions formed from the following listing of elements. Write the Lewis symbols for the monatomic ions formed from the following elements:

   (a) Cl

   (b) Na

   (c) Mg

   (d) Ca

   (e) K

   (f) Br

   (g) 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) Yard₂O

   (eastward) Li₃North

   (f) KF

4. In the Lewis structures listed here, Grand and X correspond various elements in the third menstruation 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 P_two, an unstable form of phosphorus found in high-temperature phosphorus vapor.
6. Write Lewis structures for the post-obit:

   (a) H_two

   (b) HBr

   (c) PCl_iii

   (d) SF₂

   (east) H₂CCH_two

   (f) HNNH

   (grand) H₂CNH

   (h) NO^–

   (i) N₂

   (j) CO

   (one thousand) CN^–

7. Write Lewis structures for the following:

   (a) O_ii

   (b) H₂CO

   (c) AsF₃

   (d) ClNO

   (e) SiCl_four

   (f) H_iiiO⁺

   (g) NH_four ⁺

   (h) BF₄ ⁻

   (i) HCCH

   (j) ClCN

   (k) C_ii ²⁺

8. Write Lewis structures for the following:

   (a) ClF₃

   (b) PCl₅

   (c) BF_iii

   (d) PF_vi ⁻

9. Write Lewis structures for the following:

   (a) SeF_half-dozen

   (b) XeF₄

   (c) SeCl₃ ⁺

   (d) Cl_twoBBCl_two (contains a B–B bond)

10. Write Lewis structures for:

    (a) PO₄ ³⁻

    (b) ICl₄ ⁻

    (c) SO_iii 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) SbH_iii

    (b) XeF_ii

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

13. Methanol, H₃COH, is used as the fuel in some race cars. Ethanol, C₂H₅OH, is used extensively equally motor fuel in Brazil. Both methanol and ethanol produce CO_two 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 system incorporate organic chemicals including methyl hydride (CH₄) and traces of ethylene (C₂H_iv), ethane (C_iiH₆), propyne (H_threeCCCH), and diacetylene (HCCCCH). Write the Lewis structures for each of these molecules.
15. Carbon tetrachloride was formerly used in fire extinguishers for electrical fires. Information technology 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 following electron configurations. Then, write the Lewis symbol for the common ion formed from each atom:

    (a) is ²twosouthward ⁱⁱiip ⁵

    (b) 1s ⁱⁱtwos ²2p ⁶iiis ^two

    (c) 1southward ^twoiis ⁱⁱtwop ⁶threes ²3p ^half-dozen4south ⁱⁱ3d ^x

    (d) 1due south ²2south ^twoiip ^viiiis ^twothreep ⁶4southward ²3d ¹⁰4p ⁴

    (eastward) 1due south ²twos ⁱⁱ2p ⁶3s ²3p ⁶4s ²3d ¹⁰fourp ^ane

17. The system of atoms in several biologically important molecules is given here. Complete the Lewis structures of these molecules by calculation multiple bonds and lone pairs. Do non add whatever more atoms.

    (a) the amino acid serine:

    

    (b) urea:

    

    (c) pyruvic acid:

    

    (d) uracil:

    

    (east) carbonic acid:

    

18. A chemical compound with a molar mass of about 28 chiliad/mol contains 85.7% carbon and 14.3% hydrogen by mass. Write the Lewis structure for a molecule of the compound.
19. A compound with a tooth mass of almost 42 thou/mol contains 85.7% carbon and xiv.3% hydrogen by mass. Write the Lewis structure for a molecule of the chemical compound.
20. Two arrangements of atoms are possible for a compound with a molar mass of well-nigh 45 1000/mol that contains 52.two% C, 13.1% H, and 34.7% O by mass. Write the Lewis structures for the two molecules.
21. How are unmarried, double, and triple bonds like? How practice they differ?

Glossary

double bond

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

gratis radical

molecule that contains an odd number of electrons

hypervalent molecule

molecule containing at to the lowest degree one main grouping element that has more than eight electrons in its valence trounce

Lewis construction

diagram showing lonely 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 represent each valence electron in the element or ion

lone pair

two (a pair of) valence electrons that are not used to grade a covalent bond

octet rule

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

single bail

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

triple bail

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

Solutions

Answers to Chemical science Stop of Chapter Exercises

one. (a) 8 electrons:
;

(b) eight electrons:

;

(c) no electrons

Be²⁺;

(d) eight electrons:

;

(e) no electrons

Gaⁱⁱⁱ⁺;

(f) no electrons

Li⁺;

(grand) eight electrons:

3. (a)

;

(b)

;

(c)

;

(d)

>;

(due east)

;

(f)

5.

7. (a)

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

(b)

;

(c)

;

(d)

;

(e)

;

(f)

;

(g)

;

(h)

;

(i)

;

(j)

;

(thou)

9. (a) SeF₆:
;

(b) XeF₄:

;

(c) SeCl₃ ⁺:

;

(d) Cl₂BBCl₂:

11. Two valence electrons per Pb cantlet are transferred to Cl atoms; the resulting Pb^two+ ion has a 6southward ² valence shell configuration. Two of the valence electrons in the HCl molecule are shared, and the other six are located on the Cl atom as lone pairs of electrons.

13.

fifteen.

17. (a)
;

(b)

;

(c)

;

(d)

;

(eastward)

19.

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 bail.

How To Determine The Base Element In A Covalent Bond When Drawing A Lewis Dot,

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