AP Chemistry 2.5-2.6: Lewis Diagrams, Resonance, and Formal Charge

Lewis symbol

A Lewis symbol is an element symbol surrounded by one dot for each valence electron of the atom or monatomic ion.

How do Lewis symbols represent the formation of cations and anions?

A cation is shown with fewer dots because the atom loses electrons, while an anion is shown with more dots because the atom gains electrons. Electron transfer changes which atom has the electrons but does not destroy the total number of electrons.

What does a Lewis structure depict?

A Lewis structure shows the arrangement of valence electrons, covalent bonds, and lone pairs in a molecule or polyatomic ion.

Single bond

A single bond is one shared pair of electrons between two atoms, commonly represented by a single line.

Lone pair

A lone pair is a pair of valence electrons localized on one atom and not used to form a covalent bond.

How many lone pairs and bonding pairs surround each atom in Cl₂ in its Lewis structure?

The two chlorine atoms share one bonding pair, forming a single bond. Each chlorine also has three lone pairs, giving each atom eight electrons around it.

Octet rule

The octet rule is the tendency of many main-group atoms to form bonds so that they have eight valence electrons, corresponding to a noble-gas-like configuration.

How can the typical number of covalent bonds for a second-period nonmetal be predicted from its valence electrons?

The atom generally forms enough bonds to supply the electrons needed for an octet. Thus, carbon commonly forms four bonds, nitrogen three, and oxygen two; halogens commonly form one.

Why does hydrogen not follow the octet rule?

Hydrogen is stable with two electrons in its 1s valence shell, so it follows a duet rule rather than requiring eight electrons.

Which major classes of elements commonly violate the octet rule?

Hydrogen, transition and inner-transition elements, odd-electron species, electron-deficient central atoms, and hypervalent molecules can violate the octet rule.

Double bond

A double bond consists of two shared electron pairs between two atoms and is represented by two lines.

Triple bond

A triple bond consists of three shared electron pairs between two atoms and is represented by three lines.

How do double and triple bonds help atoms satisfy the octet rule?

If single bonds and lone pairs do not give an atom an octet, a lone pair on a neighboring atom can be converted into an additional bonding pair. Two shared pairs make a double bond, and three make a triple bond.

What is the general procedure for constructing a Lewis structure?

Count total valence electrons, adjust for ionic charge, choose and draw the skeleton, connect atoms with single bonds, complete terminal-atom octets, place remaining electrons on the central atom, and form multiple bonds if needed to complete octets.

How is the total number of valence electrons adjusted for a polyatomic ion?

Add one electron for each negative charge and subtract one electron for each positive charge. The resulting total must equal the electrons represented in the completed Lewis structure.

What usually determines which atom is placed at the center of a Lewis structure?

The least electronegative atom is generally central because it is more likely to form multiple bonds. Hydrogen is almost never central, and fluorine is not placed centrally.

Why are brackets and an external charge used around a Lewis structure for a polyatomic ion?

Brackets distinguish the entire species from a neutral molecule, and the charge written outside the brackets indicates the net charge of the ion.

How many valence electrons should be counted for SiH₄, HCO₂⁻, NO⁺, and OF₂?

SiH₄ has $4+4(1)=8$ valence electrons; HCO₂⁻ has $4+1+2(6)+1=18$; NO⁺ has $5+6-1=10$; and OF₂ has $6+2(7)=20$.

What are the Lewis structures of HCN, ethane (H₃CCH₃), acetylene (HCCH), and NH₃?

HCN is $\mathrm{H-C\equiv N}$ with one lone pair on N; ethane is $\mathrm{H_3C-CH_3}$; acetylene is $\mathrm{H-C\equiv C-H}$; and ammonia is $\mathrm{NH_3}$ with one lone pair on N.

What are the Lewis structures of CO and CO₂?

CO is commonly represented as $\mathrm{C\equiv O}$ with one lone pair on each atom and formal charges $\mathrm{C^{-1}}$ and $\mathrm{O^{+1}}$. CO₂ is $\mathrm{O=C=O}$, with two lone pairs on each oxygen and no formal charges.

Free radical

A free radical is a species with an unpaired electron, usually resulting from an odd total number of valence electrons. Such a species cannot give every atom a complete octet.

How is the Lewis structure of an odd-electron molecule such as NO handled?

Count the odd number of valence electrons, place electrons to complete the more electronegative atom's octet, and use multiple bonds to give the other atom as many electrons as possible. One unpaired electron must remain.

Electron-deficient molecule

An electron-deficient molecule has a central atom with fewer than eight electrons in its Lewis structure. Examples include BeH₂ and BF₃.

Why is BF₃ usually represented with three B–F single bonds rather than a B=F double bond?

Experimental bond lengths and the observed reactivity of BF₃ support three nearly single B–F bonds with an electron-deficient boron. Boron can accept a lone pair from a donor such as NH₃.

How does NH₃ react with BF₃ in terms of Lewis structures?

The lone pair on nitrogen in NH₃ is donated to electron-deficient boron in BF₃, forming a coordinate covalent bond and giving boron a more complete valence shell.

Hypervalent molecule

A hypervalent molecule has a central atom surrounded by more than eight electrons. Expanded valence shells are possible mainly for elements in period 3 or higher, such as P in PCl₅ and S in SF₆.

How are the Lewis structures of XeF₂ and XeF₆ constructed?

In XeF₂, xenon has two Xe–F bonds and three lone pairs. In XeF₆, xenon has six Xe–F bonds and one lone pair; the extra electrons remain on the period-5 central atom.

Formal charge

Formal charge is a bookkeeping charge assigned by assuming bonding electrons are divided equally between bonded atoms; it is not the atom's actual charge in the molecule.

What equation calculates the formal charge on an atom in a Lewis structure?

Formal charge $=V-N-\frac{B}{2}$, where $V$ is the atom's valence-electron count, $N$ is its nonbonding-electron count, and $B$ is the number of bonding electrons.

What must the sum of all formal charges equal?

The formal charges in a neutral molecule must sum to zero. In a polyatomic ion, they must sum to the ion's overall charge.

What formal charges are assigned to carbon and oxygen in the common Lewis structure of CO?

For $\mathrm{C\equiv O}$ with one lone pair on each atom, carbon has formal charge $-1$ and oxygen has formal charge $+1$.

Which Lewis structure is generally preferred when comparing possible structures using formal charge?

Prefer structures with all formal charges equal to zero. If nonzero charges are unavoidable, prefer the structure with fewer and smaller formal charges, opposite charges on adjacent atoms when possible, and negative charge on the more electronegative atom.

Why is carbon-centered O=C=O the preferred Lewis structure for CO₂?

The structure with carbon in the center and two C=O double bonds gives every atom a formal charge of zero. Alternatives involving single and triple bonds or oxygen as the central atom produce less favorable formal-charge distributions.

How can formal charge help select the atom arrangement in an ion such as SCN⁻?

Formal-charge comparisons favor arrangements with the fewest atoms bearing nonzero charges and with the negative charge on the more electronegative atom. These criteria favor the arrangement with sulfur and nitrogen at the ends and carbon in the center.

What features identify the preferred Lewis structure for N₂O?

The preferred structure has oxygen at an end and nitrogen in the center, minimizes the number and magnitude of formal charges, and places any negative formal charge on oxygen rather than nitrogen.

Resonance

Resonance occurs when more than one valid Lewis structure has the same atom connectivity but different locations of multiple bonds or lone pairs. The actual species is a resonance hybrid, not a molecule rapidly switching between separate structures.

What changes and what stays the same among resonance forms?

The positions of atoms and the total number of valence electrons stay the same. Only the placement of electrons—such as π bonds, lone pairs, or formal charges—changes.

How are resonance structures generated from one another?

Keep the atomic skeleton and overall charge unchanged, then move lone pairs or multiple bonds to different locations. Atoms are not rearranged, and the total number of electrons remains constant.

What is the relationship between a resonance hybrid and its resonance forms?

The resonance hybrid is the actual electron distribution, represented collectively by all valid resonance forms. It is not a rapid alternation between separate molecules.

What bond properties commonly result from equivalent resonance forms?

Equivalent resonance forms produce equivalent bonds whose lengths and strengths are intermediate between those of single and multiple bonds because the bonding electrons are delocalized.

Why can nitrite, NO₂⁻, be represented by two resonance forms?

The N–O connectivity is fixed, but the N=O double bond can be drawn to either oxygen. The actual ion has equivalent N–O bonds with electron density delocalized between the two resonance positions.

How many equivalent resonance forms does carbonate, CO₃²⁻, have, and what do they show?

Carbonate has three equivalent resonance forms, each placing the C=O double bond on a different oxygen. The actual ion has three equivalent C–O bonds with delocalized π-electron density.

What is the relationship between Lewis structures and experimental evidence when multiple structures seem possible?

Lewis structures are models, so formal-charge rules help identify likely arrangements, while experimental observations such as bond lengths and reactivity can distinguish or support structures and reveal partial multiple-bond character.