Readings That Relate to Valence Electrons and Lewis Dot Structures

Learning Objectives

By the end of this department, y'all will be able to:

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

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

Lewis Symbols

We use Lewis symbols to describe 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:

Effigy vii.9 shows the Lewis symbols for the elements of the third catamenia of the periodic table.

Effigy 7.9 Lewis symbols illustrating the number of valence electrons for each element in the 3rd period of the periodic tabular array.

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

Likewise, they can be used to show the formation of anions from atoms, every bit shown here for chlorine and sulfur:

Figure 7.10 demonstrates the utilize of Lewis symbols to show the transfer of electrons during the formation of ionic compounds.

Figure 7.10 Cations are formed when atoms lose electrons, represented past fewer Lewis dots, whereas anions are formed by atoms gaining electrons. The total number of electrons does not change.

Lewis Structures

We also apply Lewis symbols to indicate the formation of covalent bonds, which are shown in Lewis structures, drawings that draw 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 (called alone 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 single shared pair of electrons is called a single bond. 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 element of group vii molecules (F_two, Br₂, I_ii, and At₂) form bonds like those in the chlorine molecule: one single bond between atoms and three lone pairs of electrons per atom. This allows each halogen atom to have a noble gas electron configuration. The tendency of master group atoms to form plenty bonds to obtain eight valence electrons is known as the octet dominion.

The number of bonds that an atom can form can often be predicted from the number of electrons needed to reach an octet (eight valence electrons); this is specially truthful of the nonmetals of the second menses of the periodic tabular array (C, N, O, and F). For example, each atom of a group 14 element has four electrons in its outermost beat out and therefore requires 4 more than electrons to attain an octet. These iv electrons can exist gained by forming 4 covalent bonds, as illustrated here for carbon in CCl₄ (carbon tetrachloride) and silicon in SiH_iv (silane). Considering hydrogen only needs 2 electrons to fill its valence shell, information technology is an exception to the octet rule. The transition elements and inner transition elements likewise do not follow the octet rule:

Group fifteen elements such as nitrogen have 5 valence electrons in the atomic Lewis symbol: one lone pair and three unpaired electrons. To obtain an octet, these atoms class three covalent bonds, equally in NH₃ (ammonia). Oxygen and other atoms in group xvi obtain an octet past forming two covalent bonds:

Double and Triple Bonds

As previously mentioned, when a pair of atoms shares 1 pair of electrons, we call this a single bond. However, a pair of atoms may demand to share more than one pair of electrons in social club to achieve the requisite octet. A double bond forms when two pairs of electrons are shared between a pair of atoms, equally between the carbon and oxygen atoms in CH₂O (formaldehyde) and betwixt the 2 carbon atoms in C₂H₄ (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 Dominion

For very elementary molecules and molecular ions, we can write the Lewis structures past merely pairing up the unpaired electrons on the constituent atoms. See these examples:

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

1. Decide the full number of valence (outer vanquish) electrons. For cations, decrease one electron for each positive charge. For anions, add one electron for each negative charge.
2. Draw a skeleton construction of the molecule or ion, arranging the atoms around a key atom. (By and large, the to the lowest degree electronegative element should be placed in the eye.) Connect each atom to the central atom with a single bond (one electron pair).
3. Distribute the remaining electrons as solitary pairs on the terminal atoms (except hydrogen), completing an octet around each atom.
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 SiH₄, ${\text{CHO}}_2{}^{\text{−}},$ NO⁺, and OF_ii as examples in post-obit this process:

1. Decide 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:
     

     $$\begin{array}{l}\\ {\text{SiH}}_4\\ \text{Si: 4 valence electrons/atom}\times \text{ane atom}=\mathrm{iv}\\ \underset{¯}{+\text{H: 1 valence electron/cantlet}\times \text{4 atoms}=\mathrm{iv}}\\ \\ =\text{viii valence electrons}\end{array}$$

   • For a negative ion, such equally ${\text{CHO}}_{\mathrm{ii}}{}^{\text{−}},$ we add the number of valence electrons on the atoms to the number of negative charges on the ion (ane electron is gained for each single negative charge):
     

     $$\begin{array}{l}\\ {\text{CHO}}_2{}^{\text{−}}\\ \text{C: iv valence electrons/atom}\times \text{i cantlet}=4\\ \text{H: 1 valence electron/atom}\times \text{1 cantlet}=1\\ \text{O: 6 valence electrons/atom}\times \text{2 atoms}=12\\ \underset{¯}{+\text{1 boosted electron}=1}\\ \\ =\text{18 valence electrons}\end{array}$$

   • For a positive ion, such every bit NO⁺, we add the number of valence electrons on the atoms in the ion and then subtract the number of positive charges on the ion (one electron is lost for each single positive charge) from the total number of valence electrons:
     

     $$\begin{array}{}\\ \\ {\text{NO}}^{+}\\ \text{N: 5 valence electrons/atom}\times \text{1 atom}=5\\ \\ \text{O: 6 valence electron/atom}\times \text{1 atom}=\mathrm{half\; dozen}\\ \underset{¯}{+\text{−ane electron (positive accuse)}=\mathrm{-1}}\\ \\ =\text{x valence electrons}\end{array}$$

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

     $$\begin{array}{l}\\ {\text{OF}}_{\text{2}}\\ \text{O: 6 valence electrons/atom}\times \text{i atom}=6\\ \underset{¯}{+\text{F: 7 valence electrons/atom}\times \text{two atoms}=14}\\ =\text{20 valence electrons}\end{array}$$

2. Draw a skeleton structure of the molecule or ion, arranging the atoms around a primal atom and connecting each atom to the central atom with a unmarried (one electron pair) bond. (Note that we denote ions with brackets around the construction, indicating the charge outside the brackets:)
   
   When several arrangements of atoms are possible, every bit for ${\text{CHO}}_{\mathrm{two}}{}^{\text{−}},$ we must utilize experimental evidence to choose the correct 1. In general, the less electronegative elements are more than likely to be central atoms. In ${\text{CHO}}_2{}^{\text{−}},$ the less electronegative carbon cantlet occupies the primal position with the oxygen and hydrogen atoms surrounding information technology. Other examples include P in POCl₃, Due south in And so₂, and Cl in ${\text{ClO}}_4{}^{\text{−}}.$ An exception is that hydrogen is well-nigh never a central atom. As the most electronegative element, fluorine also cannot exist a central atom.
3. Distribute the remaining electrons every bit solitary 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. Place all remaining electrons on the fundamental atom.
   • For SiH_iv, ${\text{CHO}}_2{}^{\text{−}},$ and NO⁺, at that place are no remaining electrons; we already placed all of the electrons determined in Pace one.
   • For OF₂, we had xvi electrons remaining in Step 3, and we placed 12, leaving 4 to exist placed on the central atom:
     
5. Rearrange the electrons of the outer atoms to make multiple bonds with the central atom in society to obtain octets wherever possible.

Case 7.four

Writing Lewis Structures

NASA'due south Cassini-Huygens mission detected a large cloud of toxic hydrogen cyanide (HCN) on Titan, 1 of Saturn's moons. Titan also contains ethane (H₃CCH_three), acetylene (HCCH), and ammonia (NH₃). What are the Lewis structures of these molecules?

Solution

1. Step 1. Calculate the number of valence electrons.
   HCN: (i $\times$ 1) + (4 $\times$ 1) + (v $\times$ 1) = ten
   H_iiiCCH₃: (1 $\times$ iii) + (2 $\times$ iv) + (i $\times$ 3) = 14
   HCCH: (i $\times$ 1) + (ii $\times$ 4) + (one $\times$ 1) = 10
   NH₃: (5 $\times$ i) + (3 $\times$ 1) = eight
2. Step two. Draw a skeleton and connect the atoms with single bonds. Remember that H is never a central atom:
   
3. Step three. Where needed, distribute electrons to the terminal atoms:
   
   HCN: half dozen electrons placed on Northward
   H₃CCH_iii: no electrons remain
   HCCH: no terminal atoms capable of accepting electrons
   NH_three: no terminal atoms capable of accepting electrons
4. Step 4. Where needed, identify remaining electrons on the central atom:
   
   HCN: no electrons remain
   H₃CCH_iii: no electrons remain
   HCCH: four electrons placed on carbon
   NH_iii: two electrons placed on nitrogen
5. Step 5. Where needed, rearrange electrons to course multiple bonds in social club to obtain an octet on each atom:
   HCN: class two more C–N bonds
   H_threeCCH₃: all atoms have the correct number of electrons
   HCCH: form a triple bail between the two carbon atoms
   NH_iii: all atoms have the correct number of electrons
   

Check Your Learning

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

Answer:

How Sciences Interconnect

Fullerene Chemistry

Carbon, in various forms and compounds, has been known since prehistoric times, . Soot has been used every bit a pigment (oftentimes chosen carbon black) for thousands of years. Charcoal, high in carbon content, has also been critical to homo development. Carbon is the key additive to iron in the steelmaking process, and diamonds have a unique place in both civilization and industry. With all this usage came significant study, specially with the emergence of organic chemistry. And even with all the known forms and functions of the element, scientists began to uncover the potential for even more varied and extensive carbon structures.

Every bit early on as the 1960s, chemists began to observe complex carbon structures, merely they had little evidence to support their concepts, or their piece of work did not make it into the mainstream. Eiji Osawa predicted a spherical form based on observations of a similar structure, but his piece of work was not widely known outside Japan. In a like way, the well-nigh comprehensive advance was likely computational chemist Elena Galpern's, who in 1973 predicted a highly stable, lx-carbon molecule; her work was also isolated to her native Russia. All the same subsequently, Harold Kroto, working with Canadian radio astronomers, sought to uncover the nature of long carbon chains that had been discovered in interstellar space.

Kroto sought to utilise a machine developed by Richard Smalley's team at Rice University to larn more most these structures. Together with Robert Curl, who had introduced them, and three graduate students—James Heath, Sean O'Brien, and Yuan Liu—they performed an intensive series of experiments that led to a major discovery.

In 1996, the Nobel Prize in Chemical science was awarded to Richard Smalley (Figure 7.11), Robert Curl, and Harold Kroto for their work in discovering a new class of carbon, the C_lx buckminsterfullerene molecule (Figure 7.one). An unabridged class of compounds, including spheres and tubes of various shapes, were discovered based on C_sixty. This blazon of molecule, chosen a fullerene, shows hope 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 commitment systems. They also possess unique electronic and optical properties that have been put to adept employ in solar powered devices and chemical sensors.

Figure seven.eleven 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 US Senate honored him every bit the "Father of Nanotechnology." (credit: United States Section of Energy)

Exceptions to the Octet Rule

Many covalent molecules take central atoms that do 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 have an unpaired electron.
• Electron-scarce molecules accept a central atom that has fewer electrons than needed for a element of group 0 configuration.
• Hypervalent molecules have a central atom that has more electrons than needed for a element of group 0 configuration.

Odd-electron Molecules

We call molecules that contain 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 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 minor changes:

1. Determine the total number of valence (outer beat) electrons. The sum of the valence electrons is five (from Due north) + 6 (from O) = 11. The odd number immediately tells us that nosotros have a free radical, so we know that non every cantlet can accept 8 electrons in its valence trounce.
2. Describe a skeleton construction of the molecule. We can easily draw a skeleton with an Northward–O unmarried bond:
   N–O
3. Distribute the remaining electrons as lone pairs on the concluding atoms. In this case, there is no central cantlet, then we distribute the electrons around both atoms. We give eight electrons to the more electronegative atom in these situations; thus oxygen has the filled valence shell:
   
4. Identify all remaining electrons on the fundamental atom. Since there are no remaining electrons, this stride does not employ.
5. Rearrange the electrons to make multiple bonds with the primal atom in lodge to obtain octets wherever possible. We know that an odd-electron molecule cannot have an octet for every atom, but we desire to get each cantlet as close to an octet as possible. In this case, nitrogen has only five electrons around it. To motion closer to an octet for nitrogen, nosotros take one of the lone pairs from oxygen and utilize it to form a NO double bond. (We cannot take another lonely pair of electrons on oxygen and class a triple bond considering nitrogen would and so have ix electrons:)
   

Electron-scarce Molecules

We will also run across a few molecules that contain central atoms that do not have a filled valence shell. More often than not, these are molecules with central atoms from groups ii and 13, outer atoms that are hydrogen, or other atoms that practise non form multiple bonds. For example, in the Lewis structures of beryllium dihydride, BeH_two, and boron trifluoride, BF_iii, the glucinium and boron atoms each have only 4 and half-dozen electrons, respectively. It is possible to draw a construction with a double bond between a boron atom and a fluorine cantlet in BF₃, satisfying the octet rule, only experimental testify indicates the bond lengths are closer to that expected for B–F unmarried bonds. This suggests the best Lewis construction has iii B–F single bonds and an electron deficient boron. The reactivity of the compound is as well consistent with an electron deficient boron. However, the B–F bonds are slightly shorter than what is actually expected for B–F single bonds, indicating that some double bond character is found in the actual molecule.

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

Hypervalent Molecules

Elements in the 2nd flow of the periodic table (north = two) can arrange but 8 electrons in their valence shell orbitals because they have only iv valence orbitals (ane 2due south and three 2p orbitals). Elements in the tertiary and college periods (n ≥ 3) accept more than four valence orbitals and can share more than four pairs of electrons with other atoms because they accept empty d orbitals in the aforementioned crush. Molecules formed from these elements are sometimes chosen hypervalent molecules. Effigy seven.12 shows the Lewis structures for two hypervalent molecules, PCl₅ and SF_half-dozen.

Figure vii.12 In PCl₅, the fundamental atom phosphorus shares v pairs of electrons. In SF_{half dozen}, sulfur shares six pairs of electrons.

In some hypervalent molecules, such equally IF₅ and XeF₄, some of the electrons in the outer trounce of the central cantlet are solitary pairs:

When we write the Lewis structures for these molecules, we find that we accept electrons left over later on filling the valence shells of the outer atoms with viii electrons. These boosted electrons must be assigned to the central atom.

Example 7.5

Writing Lewis Structures: Octet Rule Violations

Xenon is a noble gas, simply it forms a number of stable compounds. We examined XeF₄ before. What are the Lewis structures of XeF_ii and XeF₆?

Solution

Nosotros can draw the Lewis construction of any covalent molecule by following the six steps discussed before. In this case, we can condense the last few steps, since not all of them utilise.

1. Step 1. Calculate the number of valence electrons:
   XeF_ii: viii + (2 $\times$ 7) = 22
   XeF₆: eight + (6 $\times$ 7) = l
2. Step two. Draw a skeleton joining the atoms past single bonds. Xenon volition be the key cantlet because fluorine cannot be a cardinal cantlet:
   
3. Pace 3. Distribute the remaining electrons.
   XeF₂: We place three alone pairs of electrons around each F atom, bookkeeping for 12 electrons and giving each F atom viii electrons. Thus, 6 electrons (three lone pairs) remain. These lone pairs must be placed on the Xe cantlet. This is acceptable considering Xe atoms have empty valence trounce d orbitals and can accommodate more than 8 electrons. The Lewis structure of XeF_ii shows ii bonding pairs and three lone pairs of electrons around the Xe atom:
   
   XeF_vi: We place three alone pairs of electrons around each F atom, accounting for 36 electrons. Two electrons remain, and this lone pair is placed on the Xe cantlet:
   

Cheque Your Learning

The halogens course a course 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 ${\text{ICl}}_{\mathrm{iv}}{}^{\text{−}}.$

Answer:

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Source: https://openstax.org/books/chemistry-2e/pages/7-3-lewis-symbols-and-structures

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