Chemical Bonding, Molecular Geometry, and Hybridization

Vocabulary practice flashcards covering chemical bonding types, lattice energy, Lewis dot structures, formal charge, resonance, octet rule exceptions, VSEPR geometries, bond polarity, and orbital hybridization.

Valence Electrons

Electrons located in the outermost subshells ($s$ and $p$ subshells only) of an atom that increase across periods.

Octet Rule

The chemical principle stating that elements tend to gain, lose, or share electrons until they have $8$ valence electrons.

Ionic Bonding

Chemical bonding resulting from the electrostatic attraction between oppositely charged ions, formed by the transfer of electrons from a metal (cation) to a nonmetal (anion).

Covalent Bonding

Chemical bonding characterized by the sharing of electron pairs between nonmetal atoms to satisfy the octet rule.

Metallic Bonding

Chemical bonding found in metals where valence electrons move freely throughout a lattice of metal atoms.

Lattice Energy

The energy required to completely separate one mole of a solid ionic compound into its gaseous ions; it increases as ionic charges increase and ionic radii decrease.

Lattice Energy Trend

Latrice increases = charges of ions increases = radii decreases

Strength of Ionic Charges

bigger charges= stronger bond

closer in distance= stronger interaction

smaller ion= closer and smaller attraction

Physical Properties of Ionic Compounds

Characteristics of ionic compounds including being hard and brittle, exhibiting high melting points, dissolving in water, and conducting electricity when dissolved in water. They form solids; 3D crystal lattice.

Polar Covalent Bond

A covalent bond formed by the unequal sharing of electrons, occurring when the electronegativity difference between the bonded atoms is between $0.5$ and $2$.

Nonpolar Covalent Bond

A covalent bond formed by the equal sharing of electrons, characterized by an electronegativity difference between $0$ and $0.4$, such as in bonds between identical atoms or $C-H$ bonds.

Dipole Moment

A vector quantity established when two opposite charges are separated by a distance; it increases as bond length decreases and electronegativity difference increases.

Potential Energy Dip (Covalent Bond)

The minimum point on a potential energy versus internuclear distance curve where attractive and repulsive forces are balanced, corresponding to the bond energy and bond length.

Physical Properties of Covalent Compounds

Properties typical of covalent compounds, including low melting and boiling points, nonelectrolyte behavior in water, and formation between nonmetals (or when ionic charge would be $> 4$).

Potential Energy vs Internuclear Distance Graph Trends

small atom= dip closer to zero

more bond= deeper dip= more stable bond

Lewis Structure

A structural diagram representing covalent molecules that shows how valence electrons are shared as bonds and distributed as unshared lone pairs to satisfy the octet rule.

Lewis Structure Rules

C and N are central atoms

H, Cl, Br, F, and I are in the periphery and have one bond

Nitrogen has 3 bonds

Carbon has 4 bonds

Isomers

Molecules that share the exact same molecular formula but have different arrangements of atoms and different bonding connections.

Formal Charge

A bookkeeping charge assigned to an atom in a molecule, calculated as: $\text{Formal Charge} = \text{Valence Electrons} - \frac{1}{2}(\text{Bonding Electrons}) - \text{Nonbonding Electrons}$.

Criteria for Best Lewis Structure

The most favorable resonance structure is the one where formal charges are closest to zero and any negative formal charges are located on the most electronegative atoms.

Resonance Structures

Two or more valid Lewis structures for the same molecule or polyatomic ion that differ only in the arrangement of electrons, often involving the delocalization of double bonds.

Aromatic Molecules

Organic molecules possessing resonance-stabilized ring structures, such as benzene, typically represented by a hexagon containing a dashed or solid circle to show electron delocalization.

Hypervalent Molecules

Molecules or polyatomic ions containing a central atom that accommodates an expanded octet (more than $8$ valence electrons), commonly observed with elements such as $P$, $S$, $Cl$, $Br$, $I$, and $Xe$.

Odd Octet Exception

molecule and polyatomic ions that have an odd number of electrons

Incomplete Octet Exceptions

Stable molecules in which the central atom possesses fewer than $8$ valence electrons, specifically seen in beryllium (stable with $4$ electrons) and boron (stable with $6$ electrons).

Strength and Length of Covelant Bonds

more bonds= bond length decrease= bond enthalpy increase

smaller atomic radii= shorter bond length

Bond Order

The number of electron pairs shared between two bonded atoms; an increase in bond order causes bond length to decrease and bond enthalpy to increase.

VSEPR Model

The Valence Shell Electron Pair Repulsion model, an approach used to predict three-dimensional molecular geometries based on minimizing repulsions between electron domains around a central atom.

Electron Domain

A region around a central atom occupied by electrons; single bonds, double bonds, triple bonds, and lone pairs of electrons each count as one electron domain.

Electron-Domain Geometry vs. Molecular Geometry

Electron-domain geometry is determined by the total spatial arrangement of all electron domains, whereas molecular geometry describes the positions of the bonded atoms only.

Linear

2 electron domains

2 bonds, 0 lone pairs

angle=180

Trigonal Planar

3 bond domains

3 bonds, 0 lone pairs

angle=120

Bent

3 domains

2 bonds, 1 lone pair

angle=118

Tetrahedral

4 domains

4 bonds, 0 lone pairs

angle=109.5

Trigonal Pyramidal

4 domains

3 bonds, 1 lone pair

angle=107

Bent (2,2)

4 domains

2 bonds, 2 lone pairs

angle=105

Trigonal Bipyramidal

5 domains

5 bonds, 0 lone pairs

angle=(90,120)

seesaw

5 domains

4 bonds, 1 lone pair

angle=(88,118)

T-shaped

5 domains

3 bonds, 2 lone pairs

angle=88

Linear (2,3)

5 domains

2 bonds, 3 lone pairs

angle=180

Octahedral

6 domains

6 bonds, 0 lone pairs

angle=(90,90)

Square Pyramidal

6 domains

5 bonds, 1 lone pair

angle=(88,88)

Square Planar

6 domains

4 bonds, 2 lone pairs

angle=90

Lone Pair Effect on Bond Angles

As the number of nonbonding electron domains increases, bond angles compress and decrease because nonbonding pairs and multiple bonds have higher electron charge densities and occupy more space.

Sigma ($\sigma$) and Pi ($\pi$) Bonds

Types of covalent bonds where a single bond consists of $1\,\sigma$ bond, a double bond consists of $1\,\sigma$ and $1\,\pi$ bond, and a triple bond consists of $1\,\sigma$ and $2\,\pi$ bonds.

Electronegativity

A measure of an atom's ability to attract bonded electrons toward itself, increasing toward the top-right of the periodic table, with fluorine being the most electronegative element.

Bond Dipole

The dipole moment attributed to an individual bond, behaving as a vector quantity that has both magnitude and direction.

Dipole Moment Trend

increases when:

bond length decreases

electronegativity differences increase

Dipole Moment of 2+ atom molecule

depends on polarities of bonds and geometry of molecule

Nonpolar Symmetrical $AB_n$ Geometries

Symmetrical shapes—linear, trigonal planar, tetrahedral, square planar, trigonal bipyramidal, and octahedral—where identical bond dipoles cancel each other out, making the overall molecule nonpolar.

Determining if Polar

CH bonded with any atom is polar

OH, NH, FH is polar and does hydrogen bonding

Bent, trigonal pyramidal, seesaw, t-shaped, or square pyramidal is polar

not symmetrical is polar

Determining if Nonpolar

only CH is nonpolar

same atom bonded to itself is nonpolar

linear, trigonal planar, tetrahedral, square planar, trigonal bipyramidal, and octahedral is nonpolar

Hybridization

The theoretical process of mixing different atomic orbitals on a central atom to create equivalent hybrid orbitals that account for observed molecular shapes.

Hybrid Orbitals

Orbitals produced by mixing atomic orbitals of an atom, allowing electrons to be shared with surrounding atoms in geometries matching the VSEPR model.

$sp$ Hybridization

A hybridization state formed by mixing one $s$ and one $p$ orbital, corresponding to $2$ electron domains and a linear geometry with a $180^\circ$ bond angle.

$sp^2$ Hybridization

A hybridization state formed by mixing one $s$ and two $p$ orbitals, corresponding to $3$ electron domains and a trigonal planar geometry with a $120^\circ$ bond angle.

$sp^3$ Hybridization

A hybridization state formed by mixing one $s$ and three $p$ orbitals, corresponding to $4$ electron domains and a tetrahedral geometry with a $109.5^\circ$ bond angle.