Chemical Bonds: Types and Polarity

• Chemical bonds are attractive forces that hold atoms together forming molecules or compounds.
• These bonds dictate the structure, properties, and reactivity of matter.
• Understanding bonding and molecular structure is central to understanding chemistry.

Types of Chemical Bonds

• Ionic bonds result from the electrostatic attraction between oppositely charged ions.
• These ions are formed through the transfer of electrons between atoms.
• Typically formed between metals and nonmetals.
• Covalent bonds involve the sharing of electron pairs between atoms.
• This sharing leads to a more stable electron configuration for each atom.
• Commonly formed between nonmetals.
• Metallic bonds are found in metals, where electrons are delocalized throughout a lattice of metal atoms.
• This electron "sea" allows metals to conduct electricity and heat.

Electronegativity and Bond Polarity

• Electronegativity is a measure of an atom's ability to attract electrons in a chemical bond.
• Differences in electronegativity between bonded atoms lead to bond polarity.
• Nonpolar covalent bonds occur when electrons are shared equally (electronegativity difference is small).
• Polar covalent bonds occur when electrons are shared unequally (electronegativity difference is significant).
• The atom with higher electronegativity acquires a partial negative charge (δ-).
• The atom with lower electronegativity acquires a partial positive charge (δ+).
• Ionic bonds can be considered the extreme case of polar covalent bonds.

Lewis Structures

• Lewis structures are diagrams that show bonding between atoms of a molecule and lone pairs of electrons.
• They help visualize the electron distribution in molecules.
• The octet rule states that atoms tend to gain, lose, or share electrons to achieve a full outer shell of eight electrons.
• Hydrogen is an exception to the octet rule, requiring only two electrons to complete its valence shell.
• To draw Lewis structures:
• Determine the total number of valence electrons.
• Draw a skeletal structure with the least electronegative atom in the center (except H).
• Distribute electrons to form single bonds between atoms.
• Complete the octets of terminal atoms.
• Place any remaining electrons on the central atom.
• If the central atom lacks an octet, form multiple bonds.
• Formal charge helps determine the most plausible Lewis structure when multiple structures are possible.
• It is calculated as: Formal charge = (Valence electrons) - (Non-bonding electrons) - (1/2 Bonding electrons).
• Resonance occurs when multiple valid Lewis structures can be drawn for a molecule.
• The actual structure is a resonance hybrid, an average of the contributing resonance structures.
• Resonance structures contribute to the stability of the molecule.
• Exceptions to the octet rule exist.
• Some molecules have central atoms with fewer than eight electrons (e.g., BF3).
• Some molecules have central atoms with more than eight electrons (expanded octet, e.g., SF6).
• Odd-electron species (radicals) have an unpaired electron.

VSEPR theory

• Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the geometry of molecules based on the repulsion between electron pairs around a central atom.
• Electron pairs, both bonding and non-bonding (lone pairs), arrange themselves to minimize repulsion.
• The basic steps for applying VSEPR theory include:
• Draw the Lewis structure of the molecule.
• Count the number of electron groups (atoms bonded to the central atom and lone pairs) around the central atom.
• Determine the electron-group geometry based on the number of electron groups.
• Determine the molecular geometry based on the arrangement of bonded atoms, ignoring lone pairs.
• Electron-group geometries:
• Linear: 2 electron groups. Bond angle 180°.
• Trigonal planar: 3 electron groups. Bond angle 120°.
• Tetrahedral: 4 electron groups. Bond angle 109.5°.
• Trigonal bipyramidal: 5 electron groups. Bond angles 90°, 120°, and 180°.
• Octahedral: 6 electron groups. Bond angles 90° and 180°.
• Molecular geometries describe the arrangement of atoms in space:
• Linear: 2 bonded atoms, 0 lone pairs.
• Bent: 2 bonded atoms, 1 or 2 lone pairs on a trigonal planar or tetrahedral electron-group geometry, respectively.
• Trigonal planar: 3 bonded atoms, 0 lone pairs.
• Trigonal pyramidal: 3 bonded atoms, 1 lone pair on a tetrahedral electron-group geometry.
• Tetrahedral: 4 bonded atoms, 0 lone pairs.
• See-saw: 4 bonded atoms, 1 lone pair on a trigonal bipyramidal electron-group geometry.
• T-shaped: 3 bonded atoms, 2 lone pairs on a trigonal bipyramidal electron-group geometry.
• Linear: 2 bonded atoms, 3 lone pairs on a trigonal bipyramidal electron-group geometry.
• Square pyramidal: 5 bonded atoms, 1 lone pair on an octahedral electron-group geometry.
• Square planar: 4 bonded atoms, 2 lone pairs on an octahedral electron-group geometry.
• Lone pairs exert greater repulsive forces than bonding pairs, affecting bond angles.
• Multiple bonds are treated as single electron groups for VSEPR purposes.

Molecular Polarity

• Molecular polarity depends on both bond polarity and molecular geometry.
• A molecule is polar if it has polar bonds arranged asymmetrically, resulting in a net dipole moment.
• A molecule is nonpolar if:
• It has nonpolar bonds.
• It has polar bonds arranged symmetrically so that the bond dipoles cancel out.
• Molecular polarity influences intermolecular forces and physical properties like boiling point and solubility.

Valence Bond Theory

• Valence bond (VB) theory describes covalent bond formation in terms of overlapping atomic orbitals.
• A covalent bond forms when half-filled atomic orbitals on two atoms overlap.
• The overlapping region has increased electron density, leading to bonding.
• Hybridization is the mixing of atomic orbitals to form new hybrid orbitals with different shapes and energies, suitable for bonding.
• Common types of hybridization include:
• sp: Mixing of one s and one p orbital, resulting in two sp hybrid orbitals (linear geometry).
• sp2: Mixing of one s and two p orbitals, resulting in three sp2 hybrid orbitals (trigonal planar geometry).
• sp3: Mixing of one s and three p orbitals, resulting in four sp3 hybrid orbitals (tetrahedral geometry).
• sp3d: Mixing of one s, three p, and one d orbitals, resulting in five sp3d hybrid orbitals (trigonal bipyramidal geometry).
• sp3d2: Mixing of one s, three p, and two d orbitals, resulting in six sp3d2 hybrid orbitals (octahedral geometry).
• Sigma (σ) bonds are formed by end-to-end overlap of orbitals, with electron density concentrated along the bond axis.
• Pi (π) bonds are formed by sideways overlap of p orbitals, with electron density above and below the bond axis.
• Single bonds are sigma bonds.
• Double bonds consist of one sigma and one pi bond.
• Triple bonds consist of one sigma and two pi bonds.
• Rotation around a single (sigma) bond is relatively free.
• Rotation around double or triple bonds is restricted due to the presence of pi bonds.

Molecular Orbital Theory

• Molecular orbital (MO) theory describes the electronic structure of molecules in terms of molecular orbitals, which extend over the entire molecule.
• Atomic orbitals combine to form molecular orbitals.
• Bonding molecular orbitals are lower in energy than the original atomic orbitals, promoting bonding.
• Antibonding molecular orbitals are higher in energy than the original atomic orbitals, opposing bonding.
• The number of molecular orbitals equals the number of atomic orbitals combined.
• Electrons fill molecular orbitals according to the Aufbau principle, Hund's rule, and the Pauli exclusion principle.
• Bond order is a measure of the number of bonds between two atoms in a molecule.
• It is calculated as: Bond order = (Number of electrons in bonding MOs - Number of electrons in antibonding MOs) / 2.
• A bond order of zero indicates that a molecule is unstable and unlikely to exist.
• MO diagrams show the relative energies of molecular orbitals and the filling of those orbitals with electrons.
• MO theory can explain the magnetic properties of molecules:
• Paramagnetic molecules have unpaired electrons and are attracted to a magnetic field.
• Diamagnetic molecules have all electrons paired and are repelled by a magnetic field.
• MO theory provides a more accurate description of bonding in molecules with resonance or delocalized electrons compared to Lewis structures and valence bond theory.

Intermolecular Forces

• Intermolecular forces (IMFs) are attractive forces between molecules.
• They are weaker than intramolecular forces (chemical bonds) but influence physical properties like boiling point, melting point, and viscosity.
• Types of IMFs:
• Dipole-dipole forces: occur between polar molecules.
• Hydrogen bonding: a strong type of dipole-dipole interaction between a hydrogen atom bonded to a highly electronegative atom (N, O, or F) and a lone pair on another electronegative atom.
• London dispersion forces (also called van der Waals forces): temporary, induced dipoles due to momentary fluctuations in electron distribution. Present in all molecules, but the dominant IMF in nonpolar molecules.
• The strength of London dispersion forces increases with molecular size and surface area.
• IMFs affect physical properties:
• Higher IMFs generally lead to higher boiling points and melting points.
• Stronger IMFs result in higher viscosity and surface tension in liquids.
• Solubility is affected by IMFs. "Like dissolves like": polar substances dissolve in polar solvents, and nonpolar substances dissolve in nonpolar solvents.