Lewis Structures and Octet Rule

• Ionic and covalent bonds are important types of chemical bonds.
• Electron configurations predict the number of bonds an atom can form.

Lewis Symbols

• Lewis symbols are electron configurations used to predict bonds by representative elements.
• The Lewis dot symbol of an element consists of the chemical symbol with one or more dots around it.
• Each dot represents a valence electron.
• Fluorine has 7 electrons in its outermost principal energy level (n = 2), so it has 7 dots in its Lewis symbol.
• For representative elements, the number of valence electrons is the same as the group number.
• Representative metals form ions by losing all valence electrons.
• Nonmetal elements form negative ions by acquiring electrons until they are isoelectronic with a noble gas atom.
• Ions of nonmetals have 8 electrons in their valence shell.
• Sulfide and chloride ions have net charges of –2 and -1, respectively, because the atoms needed 2e- and 1e-, respectively, to become isoelectronic with noble gas argon.
• The Lewis symbol of an element has the element symbol surrounded by between 1 and 8 dots representing valence electrons.
• Removing two electrons (dots) from ·Ca· gives just Ca2+ as the Lewis symbol of a calcium ion because this ion has lost its valence shell electrons.

Covalent Bonding and Lewis Structures

• Lewis symbols of noble gas elements show eight electrons corresponding to filled s and p subshells.
• The outer electron configuration is related to the chemical and physical traits of an element.
• Atoms gain or lose electrons until they have the same valence electrons as noble gas atoms, that is, eight.
• The octet rule states that when forming bonds, atoms of representative elements tend to gain, lose, or share electrons until they have eight electrons in the valence shell.
• Molecules are held together by bonds resulting from the sharing of electrons between two atoms in a manner consistent with the octet rule.
• A simple covalent bond forms when two atoms in a molecule share a pair of electrons.
• In hydrogen chloride formation, the dash represents a covalent bond, or a pair of electrons shared by both the H atom and the Cl atom.
• By sharing the electron pair, the Cl atom has eight valence shell electrons.
• The stability of a bond results from both atoms acquiring a noble gas configuration.
• Hydrogen is an exception to the octet rule and only needs two electrons to achieve a filled outer energy level, becoming isoelectronic with helium.
• Electron pairs on the Cl atom not involved in bonding are called lone pairs, unshared pairs, and nonbonding electrons.
• In some cases, two or three pairs of electrons are shared by two atoms to reach an octet, forming multiple bonds.
• A double bond is a covalent bond in which two pairs of electrons are shared between two atoms.
• Atoms joined by a double bond lie closer together than atoms joined by a single bond.
• Nitrogen molecules (N2) contain a triple bond.
• Lewis structures represent the covalent bonding and location of unshared electron pairs within molecules and polyatomic ions.

Steps for writing valid Lewis Structures

• Arrange the atoms in a reasonable skeletal form, placing the unique atom in the center, and determine what atoms are bonded to each other.
• Count the valence electrons and add one electron for each unit of charge in polyatomic anions.
• Connect the central atom to the surrounding atoms with single bonds.
• First add unshared pairs to complete the octets of atoms bonded to the central atom (except for hydrogen).
• Add any remaining unshared pairs to the central atom, keeping in mind the maximum number of electrons in step 2.
• If the octet rule is satisfied for each atom, the structure is correct.
• If the octet rule is not met for the central atom, and there is a shortage of valence electrons, write double or triple bonds between the central atom and surrounding atoms.

Electronegativity

• Chemical bonds are rarely purely covalent or completely ionic, but rather exhibit both.
• Atoms exhibit varying tendencies in their ability to attract and hold free electrons in the gas phase.
• Electronegativity describes the ability of an atom within a molecule to attract a shared electron pair toward itself.
• Linus Pauling developed a method for determining the relative electronegativities of the elements.
• Electronegativity values increase from left to right across a period and decrease within a group from top to bottom.
• Differing abilities of atoms in a bond to attract the shared electron pair results in unequal sharing and a shift toward the more electronegative atom, creating a polar covalent bond.
• Pure covalent bonding occurs only in homonuclear diatomic molecules (H2, N2, Cl2), where the electronegativity difference (ΔX) is zero.
• Bonds with ΔX < 2.0 are classified as polar covalent bonds, while those with ΔX ≥ 2.0 are mostly ionic.
• Sometimes the term "percent ionic character" describes the polar nature of a bond.
• Bonds between Group 1A or 2A metals and the halogens are classified as ionic due to large electronegativity differences.
• As the electronegativity difference increases, the bond becomes more polar, and its ionic character increases.
• Electronegativity of hydrogen is similar to nonmetal elements like boron and carbon, so bonds of H to nonmetal atoms are polar covalent rather than ionic.

Formal Charge

• Formal charge helps choose the most plausible Lewis structure.
• The formal charge of an atom is what the charge of an atom seems to have in a Lewis structure.
• All nonbonding electrons belong entirely to the atom in which they are found, while bonding electrons are divided equally between the bonded atoms.
• The formal charge of an atom is the number of valence electrons in an isolated atom minus the number of electrons assigned to that atom in a molecule.
• A Lewis structure with no formal charges is preferred.
• The sum of formal charges must be zero in a neutral molecule and equal to the charge of the ion in a polyatomic ion.
• The most plausible Lewis structures will have negative formal charges on the more electronegative atoms.

Resonance Structures

• Resonance structures emerge when a satisfactory electron dot structure for a molecule or polyatomic ion cannot be drawn.
• Resonance structures show the correct number of valence electrons and satisfy the octet rule.
• In resonance structures, -1 charge belongs to the entire nitrite ion, and not to just one atom in the structure.
• The structure does not accurately represent what is known about the bond lengths of the N-O bonds in NO2 because the bond lengths are known to be the same, however according to the structure the double bond is shorter than the single bond.
• A composite of these structures can describe a nitrite ion.
• Contributing/resonance structure is each of the structures that contributes to the composite structure.
• The symbol, ↔ indicates resonance structures.
• In resonance, molecules are a composite or average of the two structures, with bonds intermediate between single and double bonds.
• Resonance means the need for two or more Lewis symbols to represent a particular molecule.
• Resonance structures have the same placement of atoms but different positioning of electron pairs.
• Actual bonding in N2O is a composite of three structures.

Exceptions to the Octet Rule

• Lewis structures can be drawn for many compounds by employing the the octet rule, however, structures of some compounds do not follow the rule.

Common exceptions include:

• Molecules with an incomplete octet
• Molecules with an odd number of electrons.
• Molecules in which the central atom has an expanded octet.
• Boron halides (BX3) are examples of molecules with an incomplete octet, where the boron atom has only six valence electrons.
• Two common oxides of nitrogen, NO and NO2, have odd numbers of electrons and cannot satisfy the octet rule.
• Molecules exhibiting an expanded octet require nonmetal atoms from the third period or beyond.
• Second-period elements never exceed the octet rule, while third-period elements are just as likely to do so.
• Unlike PCl3, gaseous PCl5 has a phosphorus atom joined by single bonds to five chlorine atoms, giving the phosphorus atom 10 electrons in its valence shell.
• Central atoms in SF4, SF6, ClF5, BrF5, and IF7 exhibit expanded octets. If the central atom is from the third period or beyond, complete the octets of the surrounding atoms first, and then complete the central atom. If extra electron pairs remain, place them on the central atom.
• Exceptions to the octet rule indicate that a completed octet might not be necessary for covalent bonding.
• Within the Lewis framework, sharing electron pairs leads to the covalent bond, with shared pairs acting to attract both atoms.

Explore Lewis symbols, the octet rule, and ion formation. Understand how atoms achieve stability by gaining, losing, or sharing electrons. Covers valence electrons and isoelectronic species.