Lewis Structures Study Pack

Valence Electrons and Their Source

• •Valence electrons are the electrons in an atom's outermost principal energy level; they are the only electrons involved in chemical bonding.
• •For main-group elements, the group number in the periodic table directly gives the number of valence electrons (e.g., oxygen in Group 16 has 6 valence electrons).
• •Transition metals are treated differently and their Lewis dot representations are less commonly used because d-electron involvement in bonding is more complex.

Lewis Dot Symbols for Individual Atoms

• •A Lewis dot symbol consists of the element's chemical symbol surrounded by dots, one dot for each valence electron.
• •Dots are placed on the four sides of the symbol (top, bottom, left, right) one at a time before pairing; this convention reflects the four available orbitals in the s and p sublevels.
• •Paired dots on one side represent a lone pair (a set of two non-bonding electrons), while single dots represent unpaired electrons available to form bonds.

The Octet Rule

• •Main-group elements from Period 2 (carbon, nitrogen, oxygen, fluorine) almost always obey the octet rule strictly, ending up with exactly eight electrons around them.
• •When two atoms share a pair of electrons, both atoms count those shared electrons toward their own octets, which is why covalent bonding satisfies the octet rule for both partners simultaneously.

Hydrogen and the Duet Rule

• •Hydrogen only needs two electrons to fill its 1s orbital, so it follows a duet rule rather than the octet rule.
• •Helium is the noble gas analogue for hydrogen, and any Lewis structure showing hydrogen with more than two electrons is incorrect.

Expanded Octets in Period 3 and Beyond

• •Elements such as sulfur (S), phosphorus (P), chlorine (Cl), and xenon (Xe) can have more than eight electrons around the central atom because their third (or higher) principal energy level includes d orbitals that can participate in bonding.
• •SF₆ is a classic example: sulfur is surrounded by 12 bonding electrons (six single bonds to fluorine), exceeding the octet.

Electron-Deficient Molecules

• •Some molecules, notably those containing boron or beryllium as the central atom, are stable with fewer than eight electrons around the central atom.
• •BF₃ has only six electrons around boron and is electron-deficient, making it a strong Lewis acid that readily accepts an electron pair from a donor.

• •Step 1 — Count Total Valence Electrons
• •Sum the valence electrons contributed by every atom in the molecule; for polyatomic ions, add one electron per negative charge or subtract one per positive charge.
• •For example, NO₂⁻ has 5 (N) + 6×2 (O) + 1 (negative charge) = 18 total valence electrons to place.
• •Step 2 — Assemble the Skeleton Structure
• •Identify the central atom — usually the least electronegative atom that is not hydrogen (hydrogen is always terminal).
• •Connect the central atom to each surrounding atom with a single bond, and subtract two electrons from the total count for each bond drawn.
• •Step 3 — Distribute Remaining Electrons as Lone Pairs
• •Place lone pairs on terminal atoms first, satisfying their octets (or duets for hydrogen) before addressing the central atom.
• •Place any leftover electrons as lone pairs on the central atom.
• •Step 4 — Form Multiple Bonds If Needed
• •If the central atom still has fewer than eight electrons after Step 3, convert one lone pair from an adjacent terminal atom into a shared bonding pair, creating a double bond.
• •Repeat this conversion (creating a triple bond if necessary) until the central atom satisfies the octet rule.
• •Each conversion reduces lone pairs by one and increases the bond order by one without changing the total electron count.

Single Bonds

• •A single bond consists of one shared pair of electrons (two electrons total) and is represented by a single line between atoms.
• •Single bonds are the longest and weakest of the three bond types because the atoms share minimal electron density between their nuclei.

Double Bonds

• •A double bond involves two shared pairs of electrons (four electrons) and is drawn as two parallel lines between atoms.
• •Double bonds are shorter and stronger than single bonds; for example, the C=O bond in formaldehyde (H₂CO) is shorter than a C–O single bond in an alcohol.

Triple Bonds

• •A triple bond consists of three shared pairs (six electrons) and is the shortest and strongest covalent bond type.
• •N₂ contains a triple bond (N≡N), which accounts for its exceptional stability and chemical inertness under normal conditions.