General Chemistry: Molecular Orbital Theory, Covalent Bonding, and Lewis Structures

Flashcards covering key terms, rules, orbital energies, and molecular diagrams from General Chemistry Week 4, Class 8.

Molecular Orbitals (MOs)

Orbitals in molecules where electrons exist, which are analogous to atomic orbitals and possess quantized energy levels.

Molecular Orbital Formation Conditions

The requirement that atomic orbitals must combine by adding and subtracting orbitals of the same symmetry and similar energy.

Orbital Conservation Principle

The principle stating that the total number of molecular orbitals produced must equal the total number of atomic orbitals combined.

Sigma ($\sigma$) Molecular Orbitals

Molecular orbitals formed by head-on overlap of atomic orbitals along the bond axis, such as $s-s$, $s-p$, or $p-p$ overlap.

Covalent Bonding

A chemical bonding interaction based on the sharing of valence electrons between two or more atoms.

Lewis Dot Structure Guidelines for Bonds

Representations of shared pairs of electrons using lines: single bond ($1$ line, $2$ electrons), double bond ($2$ lines, $4$ electrons), and triple bond ($3$ lines, $6$ electrons).

Octet Rule

The tendency of most atoms to form bonds until surrounded by a full valence shell: $2$ electrons for Row 1 elements, and $8$ electrons for Row 2 and higher.

Row 2 Octet Rule Exceptions

Beryllium ($Be$) and Boron ($B$), which are Row 2 elements that can exist stably with fewer than $8$ valence electrons.

Expanded Octet

A valence state occurring in elements from Row 3 and higher where central atoms utilize $d$ orbitals to hold more than $8$ valence electrons.

Formal Charge Formula

The mathematical formula defined as: $\text{Formal Charge} = (\text{valence } e^-) - (\text{lone pair } e^-) - \frac{1}{2}(\text{shared } e^-)$.

Formal Charge Guidelines for Optimal Lewis Structures

Rules stating that the best structure has the lowest overall formal charges, placing negative formal charges on the most electronegative elements and positive formal charges on the least electronegative elements.

Resonance Structures

Two or more equivalent Lewis structures drawn for a molecule where the positions of atoms remain fixed while only electron and bond locations differ.

Resonance Hybrid

The actual physical molecule in a resonance system, representing the weighted average of all equivalent resonance structures.

Polyatomic Ion

A charged species composed of multiple atoms covalently bound together whose individual atomic formal charges do not cancel out.

Phosphate ($PO_4^{3-}$)

A biologically important polyatomic ion composed of one central phosphorus atom bonded to four oxygen atoms carrying an overall net charge of $-3$.

Atomic Orbital Energy of Hydrogen ($1s$)

The energy value of the $1s$ atomic orbital of Hydrogen, which is $-13.6\,eV$.

Atomic Orbital Energy of Fluorine ($2p$)

The energy value of the $2p$ atomic orbitals of Fluorine, which is $-18.6\,eV$.

Atomic Orbital Energy of Fluorine ($2s$)

The energy value of the $2s$ atomic orbital of Fluorine, which is $-40.2\,eV$.

Molecular Orbital Diagram for Heteronuclear $HF$

An MO diagram showing the interaction between Hydrogen's $1s$ orbital ($-13.6\,eV$) and Fluorine's $2p_z$ orbital ($-18.6\,eV$) to form $\sigma_{(1s+2p_z)}^b$ and $\sigma_{(1s-2p_z)}^*$, while Fluorine's $2s$ ($-40.2\,eV$) and $2p_x, 2p_y$ remain non-bonding orbitals.

Molecular Orbital Diagram for Homonuclear $O_2$

An MO diagram illustrating the combination of valence $2s$ and $2p$ orbitals from two Oxygen atoms, filling $\sigma_{2s}$, $\sigma_{2s}^*$, $\sigma_{2p_z}$, $\pi_{2p_y}$, $\pi_{2p_x}$, and two unpaired electrons in $\pi_{2p_y}^*$ and $\pi_{2p_x}^*$ orbitals.