Types of Chemical Bonding: Covalent, Polar, Ionic, and Metallic

Covalent Bonding

• Definition: covalent bonds occur when two nonmetals (typically elements from the right side of the periodic table) share electrons.
• Octet stability: atoms tend to become more stable with eight valence electrons (the octet rule). This can be achieved by sharing electrons in the bond.
• Example setup: iodine and chlorine with their valence electrons form a shared electron pair in the middle, resulting in a covalent bond.
• Bond representation: a covalent bond is depicted by a line in structural drawings; single line for a single bond, two lines for a double bond, three lines for a triple bond.
• Concept of potential energy in bond formation:
  • When two atoms are far apart and not bonding, they are in a higher potential energy state.
  • When they come closer and form a bond, they reach a lower potential energy state (the lowest available for that arrangement).
• Polar vs nonpolar covalent bonds:
  • Nonpolar covalent bond: electrons are shared almost equally between two nonmetals.
  • Polar covalent bond: electrons are shared unequally, with one atom hogging electrons more than the other.
• Visual/metaphor for sharing: two nonmetals sharing electrons can be imagined as sharing evenly (a good sharing scenario) or one atom hogging electrons (a less fair sharing scenario).
• Electronegativity as the basis for polarity (introduced as the way to determine which atom hogs the electrons):
  • Electronegativity values indicate how strongly an atom attracts electrons.
  • The difference in electronegativity between the two atoms in a bond determines polarity.
  • If the electronegativity difference is small, the bond is nonpolar or nearly nonpolar; if it’s larger, the bond becomes polar covalent.

Electronegativity and Bond Polarity

• Key statement from unit on polarity: the difference in electronegativity between bonded atoms predicts bond type.
• Example: iodine and sulfur
  • EN(I) ≈ $2.5$, EN(S) ≈ $2.5$ → ΔEN ≈ $0$, which is a perfectly nonpolar covalent bond (electrons shared equally).
• Rule of thumb for covalent bonds (from the lesson):
  • If the difference in electronegativity ΔEN < $0.5$ → nonpolar covalent bond.
  • If ΔEN ≥ $0.5$ → polar covalent bond (the atom with higher electronegativity hogs electrons more).
• Consequences of polarity:
  • The side with the higher electronegativity becomes slightly negatively charged (partial negative, noted as a partial charge).
  • The other side becomes slightly positively charged (partial positive).
• Examples to illustrate differences:
  • Boron and chlorine:
  • EN(B) ≈ $2.0$, EN(Cl) ≈ $3.0$ → ΔEN = $1.0$ → polar covalent bond; chlorine hogs electrons more (higher EN).
  • Carbon and hydrogen:
  • EN(C) ≈ $2.5$, EN(H) ≈ $2.1$ → ΔEN = $0.4$ → nonpolar covalent (nearly evenly shared, perhaps a 60/40 arrangement rather than 50/50, but still considered nonpolar under the rule).
• Metal-nonmetal example (ionic bond):
  • Sodium and bromine:
  • Sodium is a metal (not listed on the nonmetal EN chart here) and bromine is a nonmetal; the bond is ionic because a metal tends to transfer electrons to the nonmetal.

Ionic Bonding

• Definition: ionic bonds form between metals and nonmetals (e.g., left side metals and right side nonmetals) where electrons are transferred rather than shared.
• Key behavior: the metal donates one or more electrons to the nonmetal, resulting in oppositely charged ions that attract each other.
• Conceptual distinction from covalent bonds: instead of sharing a pair of electrons, there is a transfer of electrons from the metal to the nonmetal.

Metallic Bonding

• What it is: bonding in metals and metallic alloys where electrons are not fixed to a particular nucleus; they are delocalized and can move relatively freely.
• Visual picture: a lattice of fixed metal nuclei surrounded by a sea of freely moving (delocalized) electrons.
• Why it matters: the delocalized electrons can flow, which explains why metals and metallic alloys conduct electricity well (electricity is the movement of electrons through a conductor).
• Practical consequence: this mobility of electrons contributes to properties like electrical conductivity, malleability, and ductility of metals.

Quick Reference and Connections

• Bond types by atom class:
  • Covalent bonds: typically between two nonmetals (right side of the periodic table).
  • Ionic bonds: typically between a metal and a nonmetal (left vs right sides of the table).
  • Metallic bonds: within metals and their alloys.
• Visual cues in drawings:
  • A line in a structure indicates a covalent bond (single, double, or triple depending on the number of lines).
• Octet rule and stability:
  • Atoms tend to achieve eight valence electrons for stability; covalent bonding helps achieve that through shared electrons. The term “octet” refers to this stable configuration.
• Energetics of bond formation:
  • Bond formation lowers the system’s potential energy, moving from a higher energy state (far apart) to a lower energy state (in bond).
• Electronegativity as the determinant of polarity:
  • Larger differences in electronegativity lead to more polar bonds; smaller differences lead to nonpolar bonds.
• Example recap with numbers:
  • Iodine (EN ≈ $2.5$) and sulfur (EN ≈ $2.5$): ΔEN ≈ $0$ → nonpolar covalent.
  • Boron (EN ≈ $2.0$) and chlorine (EN ≈ $3.0$): ΔEN = $1.0$ → polar covalent; chlorine hogs electrons more.
  • Carbon (EN ≈ $2.5$) and hydrogen (EN ≈ $2.1$): ΔEN = $0.4$ → nonpolar covalent (near 60/40 sharing possible).
  • Sodium (metal) and bromine (nonmetal): ionic bond (electron transfer).