Chemical Bonds: Ionic, Covalent, and Metallic

• Chemical bonds are forces of attraction that hold atoms together, thus forming molecules and compounds.

Types of Chemical Bonds

• Ionic bonds entail the transfer of electrons between atoms, usually between a metal and a nonmetal.
• Covalent bonds entail the sharing of electrons between atoms, usually between two nonmetals.
• Metallic bonds exist in metals and entail electron delocalization within a "sea" of electrons.

Ionic Bonds

• Formed through electrostatic attraction between oppositely charged ions.
• Typically occur between elements with large differences in electronegativity.
• Metals tend to lose electrons, forming positive ions (cations).
• Nonmetals tend to gain electrons, forming negative ions (anions).
• The resulting ions have stable electron configurations, often isoelectronic with noble gases.
• Examples include NaCl (sodium chloride) and MgO (magnesium oxide).
• Ionic compounds form crystal lattices, which are regular, repeating arrangements of ions.
• Lattices result in high melting and boiling points because of the strong electrostatic forces.
• Ionic compounds are typically hard and brittle.
• They conduct electricity when dissolved in water or melted, as the ions are then free to move.

Covalent Bonds

• Formed by the sharing of one or more pairs of electrons between two atoms.
• Typically occur between nonmetal atoms.
• Sharing electrons allows atoms to achieve a stable electron configuration.
• Single bond: sharing of one electron pair.
• Double bond: sharing of two electron pairs.
• Triple bond: sharing of three electron pairs.
• Bond length decreases as the number of shared electron pairs increases.
• Bond strength increases as the number of shared electron pairs increases.
• Covalent bonds can be polar or nonpolar, depending on the electronegativity difference between the atoms.
• Nonpolar covalent bond: electrons are shared equally (electronegativity difference is small or zero); examples: H2, Cl2.
• Polar covalent bond: electrons are shared unequally (electronegativity difference is significant); examples: H2O, HCl.
• The more electronegative atom in a polar bond has a partial negative charge (δ-).
• The less electronegative atom has a partial positive charge (δ+).
• Polarity results in dipole moments, which measure the separation of charge in a molecule.

Metallic Bonds

• Found in metals
• Valence electrons are delocalized and not associated with a single atom.
• This creates a "sea" of electrons surrounding positively charged metal ions.
• The electron sea allows metals to efficiently conduct electricity and heat.
• Explains metals' malleability and ductility, as the atoms can slide past each other without breaking bonds.
• Metallic bonds typically result in high melting and boiling points.
• The strength of metallic bonds depends on the number of valence electrons and the charge of the metal ions.

Electronegativity

• Measures an atom's ability to attract electrons in a chemical bond.
• Increases across a period (from left to right) and decreases down a group in the periodic table.
• Fluorine is the most electronegative element.
• Elements with large electronegativity differences tend to form ionic bonds.
• Elements with small electronegativity differences tend to form covalent bonds.

Bond Polarity and Dipole Moments

• Bond polarity arises when electrons are shared unequally in a covalent bond.
• A dipole moment (μ) measures the polarity of a bond or molecule.
• Defined as the product of the magnitude of the partial charges (δ) and the distance (d) between them: μ = δd.
• A dipole moment is a vector quantity, possessing both magnitude and direction.
• The direction points from the positive to the negative end of the polar bond.
• Molecular polarity depends on both the polarity of individual bonds and the molecular geometry.
• If bond dipoles cancel each other out due to symmetry, the molecule is nonpolar (e.g., CO2).
• If bond dipoles do not cancel each other out, the molecule is polar (e.g., H2O).

Molecular Geometry

• The three-dimensional arrangement of atoms in a molecule.
• Affects molecules' physical and chemical properties.
• Determined by the valence shell electron pair repulsion (VSEPR) theory.
• VSEPR theory states that electron pairs around a central atom arrange themselves to minimize repulsion.
• Electron pairs can be bonding pairs (shared in covalent bonds) or lone pairs (non-bonding).
• Lone pairs exert more repulsive force than bonding pairs.
• Common molecular geometries include:
  • Linear (e.g., CO2)
  • Trigonal planar (e.g., BF3)
  • Bent (e.g., H2O)
  • Tetrahedral (e.g., CH4)
  • Trigonal pyramidal (e.g., NH3)
  • Trigonal bipyramidal (e.g., PCl5)
  • Octahedral (e.g., SF6)

Valence Bond Theory

• Describes covalent bond formation as the overlap of atomic orbitals.
• Atomic orbitals combine to form sigma (σ) and pi (π) bonds.
• Sigma bonds are formed by end-to-end overlap of orbitals, with electron density concentrated along the bond axis.
• Pi bonds are formed by side-by-side overlap of p orbitals, with electron density above and below the bond axis.
• Single bonds are sigma bonds.
• Double bonds consist of one sigma bond and one pi bond.
• Triple bonds consist of one sigma bond and two pi bonds.

Hybridization

• The process of mixing atomic orbitals to form new hybrid orbitals with different shapes and energies.
• Hybridization allows atoms to form stronger and more directional bonds.
• The number of hybrid orbitals formed is equal to the number of atomic orbitals mixed.
• Common types of hybridization:
  • sp hybridization: one s orbital and one p orbital mix to form two sp hybrid orbitals (e.g., BeCl2, linear geometry).
  • sp2 hybridization: one s orbital and two p orbitals mix to form three sp2 hybrid orbitals (e.g., BF3, trigonal planar geometry).
  • sp3 hybridization: one s orbital and three p orbitals mix to form four sp3 hybrid orbitals (e.g., CH4, tetrahedral geometry).
  • sp3d hybridization: one s orbital, three p orbitals, and one d orbital mix to form five sp3d hybrid orbitals (e.g., PCl5, trigonal bipyramidal geometry).
  • sp3d2 hybridization: one s orbital, three p orbitals, and two d orbitals mix to form six sp3d2 hybrid orbitals (e.g., SF6, octahedral geometry).

Molecular Orbital (MO) Theory

• A more advanced theory that describes bonding in terms of molecular orbitals, which are formed by the combination of atomic orbitals from all atoms in the molecule.
• Atomic orbitals combine to form bonding and antibonding molecular orbitals.
• Bonding orbitals are lower in energy than the original atomic orbitals and promote bonding.
• Antibonding orbitals are higher in energy than the original atomic orbitals and weaken bonding.
• Electrons fill molecular orbitals according to the Aufbau principle, Hund's rule, and the Pauli exclusion principle.
• Bond order is calculated by subtracting the number of electrons in antibonding orbitals from the number of electrons in bonding orbitals, then dividing by two.
• A positive bond order indicates that the molecule is stable, while a zero or negative bond order indicates that the molecule is unstable.

Explore the fundamental forces that hold atoms together, forming molecules and compounds. Learn about ionic bonds through electron transfer, covalent bonds through electron sharing and metallic bonds through electron delocalization. Examples include: NaCl (sodium chloride) and MgO (magnesium oxide).