Molecular structure and bonding - Lewis concept and atomic bonding rules

Oct 26, 2026

Molecular structure and bonding explains how to form and analyze molecular compounds using key prerequisites: periodic table groups and periods, electron counting and valence electrons, electronic structures (ground and excited states), and the Lewis concept for bonding. It covers quantum numbers and atomic orbitals, shell and subshell energy levels, limits on shells and electrons, and the 2-electron requirement for a bond. Examples compare maximum bond counts in elements across the 2nd and 3rd periods, including promotion energy considerations, octet rule and hypervalence, resonance and formal charges.

LanguageEnglish

CategoryResearch & Report

FormatPDF

Pages9

Reading Time23 min

AccessFree to download

Molecular prerequisites for bonding

Periodic table, electron counting, valence

Electronic structure: ground vs excited states

Atomic shells and quantum states

Energy levels and quantum numbers

Pauli principle and Hund’s rule

Bond capacity and main-group trends

Maximum shells/subshells/orbitals and electrons

Number of bonds and main-group relation

Examples and rules in practice

Second-period bond limits and octet rule

Third-period bond expansion via d orbitals

Lewis formulas, resonance, hypervalence, and formal charges

Molecular structure and bonding To understand the formation and structure of molecular compounds, first one has to learn, recognize, use, count, take into account: • the periodic table with groups and periods, • the number of electrons and valence electrons (i.e. count electrons),(2 (K), 8 (L) = 2 + 6, 18 (M) = 2 + 6 + 10, 32 (N) = 2 + 6 + 10 + 14, etc. ), • the electronic structure (ground state, exited state),• the need of 2 electrons for a bond, •atomic orbitals and quantum states (with quantum numbers n, m, l, s),• the energy levels of the different shells, subshells, orbitals, • the maximum number of shells, subshells, orbitals, and electrons, • the numbers and/or maximum numbers of bonds, an atom likes to or can build or have, • the Lewis concept/structure. Energies of the atomic shells K, L, M, N and the subshells/orbitals s, p, d, f ← Start of 3d elements (Sc) Atomic number Z Energy increases with increasing n, and is high between the shells (except d and f) . n → size, l → shape, m → orientation, s → spin of e, Pauli principle (not all 4 quantum numbers can be equal), Hund‘s rule (spin maximizing) . It is useful to learn also the maximum number of shells (n), subshells/orbitals (s, p, d, f), and electrons an atom can have. Shell n Orbital Nunber of Number of electrons in orbitals orbital shell (2n2) The next step is to find the numbers and/or maximum numbers of bonds, an atom likes to or can have, e.g. in hydrogen compounds of group 4 to 8 elements. Main group Electron * configuration Number of possible bonds Existing H compouns Lewis formula /4 * none use – instead of : * Promotion/exitation of an 1s electron ofC, Si to 1p needs a lot of energy (460 kJ/mol for C), but is compensated by the formation of 2 additional bonds (CH4 instead of CH2) . For promotion to a d level, the energy is to high. Electron configuration and number of bonds for elements of 2. period Atom or ion Number of bonds Number * of electrons Examples Electron configuration Elements of 2. period can build not more than 4 bonds, because there are only 4 orbitals available → octett rule. Consider ground states (Pauli, Hund), exited states. * Number of electrons in bonds Electron configuration and number of bonds for elements of 3. period Atom or ion Number Number * Examples of bonds of electrons Electron configuration For n > 2 one can have more than 4 bonds, because there are empty low-level d orbitals avaiable. For elements of 3. period, the maximum number of bonds is 7. * Number of electrons in bonds Highest number of bonds is equivalent to the main group number. Element Group Number of bonds P 5 3, 5 S 6 2, 4, 6 Cl 7 1, 3, 5, 7 Some examples: SiF62-, PF5, SF6 do exist CF62-, NF5, OF6 do not exist Examples: Lewis structure of N2 , NH3 , NO2-, NO3-, S2 , SO2 , SO3 , SO32 . Those which are not necessarily resonance structures/hybrids are marked with an asterisk. → Octett rule, hypervalence, formal charges -are used instead of : * No other resonance structure

FAQs

What foundational knowledge is needed to understand molecular compound formation and structure?

You need periodic table information (groups and periods), electron and valence electron counting, electronic structure (ground/excited states), the 2-electron requirement for a bond, atomic orbitals and quantum numbers, and the Lewis concept/structure.

How do quantum numbers relate to an electron’s properties in orbitals?

The document links n to size, l to shape, m to orientation, and s to spin, and uses the Pauli principle and Hund’s rule to account for how electrons fill quantum states.

How can elements form more than four bonds, and what limits apply?

For n > 2, empty low-level d orbitals allow more than four bonds; for the 3rd period the maximum bond number is stated as 7. The 2nd period is limited to 4 bonds because only four orbitals are available, tying this to the octet rule.

Oct 26, 2026

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