Covalent Bonds

Fundamentals of Covalent Bonding

• Definition of Covalent Bonds:

  • A covalent bond is a chemical link formed between atoms that share their valence electrons.
  • Covalent bonds occur because the participating atoms have similar electronegativity values, preventing the complete transfer of electrons from one atom to another.
  • Covalent bonding typically occurs between two or more non-metals, or between non-metals and metalloids.

• Mechanism and Electron Configuration:

  • A covalent bond represents the mutual sharing of one or more pairs of valence electrons between atoms.
  • Octet Rule Stability: By sharing valence electrons, each bonded atom can fill its outer valence shell and achieve stability equivalent to a noble gas octet ($8$ valence electrons, or $2$ in the case of hydrogen).

Physical and Chemical Properties of Covalent Compounds

• State of Matter at Room Temperature:

  • Many covalent molecular substances exist as gases or liquids at standard room temperature.
  • Common representative examples include carbon dioxide ($CO_2$), water ($H_2O$), and ammonia ($NH_3$).

• Mechanical and Surface Properties of Solids:

  • Covalent compounds that form solids are characteristically:
  • Soft in texture.
  • Brittle when subjected to mechanical stress.
  • Non-lustrous (dull in appearance, lacking metallic sheen).

• Conductivity:

  • Covalent compounds do not conduct heat or electricity effectively.
  • Cause: Conductivity requires freely moving charged carriers; covalent molecules lack mobile ions or free electrons to carry an electric current or transfer thermal energy.

• Thermal Properties:

  • Covalent molecular substances typically possess low to moderate melting points due to the relatively weak intermolecular forces holding discrete molecules together compared to ionic lattice networks.

Solid State Structures: Crystalline vs. Amorphous

• Structural Arrangements of Covalent Solids:
  • Solid covalent materials exhibit two distinct structural organizations:
  • Crystal Lattice Structure (Crystalline): Displays a highly organized, repeating, long-range geometric pattern.
  • Amorphous Structure: Lacks long-range order; particles are arranged randomly without a repeating structural lattice.

Electronegativity and Bond Classification

• Electronegativity Principles:

  • Electronegativity measures the ability of an atom in a chemical bond to attract shared electrons toward itself.
  • Electron sharing between two atoms is not always equal; the difference in electronegativity values ($\Delta EN$) between the bonded elements dictates the bond type and degree of polarity.
  • A greater electronegativity difference corresponds to a more polar bond.

• Solubility Behavior Based on Polarity:

  • Polar Covalent Molecules: Dissolve in polar solvents such as water ($H_2O$).
  • Nonpolar Covalent Molecules: Insoluble in water.

• Bond Classification Thresholds:

  • Nonpolar Covalent Bond:
  • Electronegativity difference: $\Delta EN \le 0.4$
  • Formed between two or more non-metals with equal or nearly equal sharing of bonding electrons.
  • Results in an even electron density distribution with no localized partial charges on either atom (e.g., homonuclear diatomic molecules like $Cl_2$).
  • Polar Covalent Bond:
  • Electronegativity difference: $\Delta EN \ge 0.5$ (up to the ionic threshold)
  • Bonding electrons are shared unequally between the two atoms, being pulled closer to the more electronegative atom.
  • Results in an asymmetric electron cloud and the formation of partial charges: a partial negative charge ($\delta-$) on the more electronegative atom and a partial positive charge ($\delta+$) on the less electronegative atom.
  • Ionic Bond:
  • Electronegativity difference: $\Delta EN > 1.7$
  • Involves the outright transfer of electrons from a metal/cation to a non-metal/anion, forming distinct ions.

• Bond Polarity Calculations and Comparative Examples:

  • Comparison of $C-H$ vs. $N-H$ Polarity:
  • Carbon-Hydrogen bond ($C-H$): $2.5 - 2.1 = 0.4$ (Nonpolar covalent bond, $\Delta EN \le 0.4$)
  • Nitrogen-Hydrogen bond ($N-H$): $3.0 - 2.1 = 0.9$ (Polar covalent bond, $\Delta EN \ge 0.5$)
  • Conclusion: An $N-H$ bond is significantly more polar than a $C-H$ bond because it exhibits a higher electronegativity difference.
  • Phosphorus-Selenium Bond ($P-Se$):
  • Electronegativity calculation: $2.4 - 2.1 = 0.3$
  • Classification: Nonpolar covalent ($\Delta EN \le 0.4$)
  • Hydrogen-Chlorine Bond ($H-Cl$):
  • Electronegativity calculation: $3.0 - 2.1 = 0.9$
  • Classification: Polar covalent ($\Delta EN \ge 0.5$)

• Selected Reference Electronegativity Values:

  • Hydrogen ($H$): $2.1$
  • Group 1: Lithium ($Li$): $1.0$, Sodium ($Na$): $0.9$, Potassium ($K$): $0.8$, Rubidium ($Rb$): $0.8$, Cesium ($Cs$): $0.7$, Francium ($Fr$): $0.7$
  • Group 2: Beryllium ($Be$): $1.5$, Magnesium ($Mg$): $1.2$, Calcium ($Ca$): $1.0$, Strontium ($Sr$): $0.9$, Barium ($Ba$): $0.9$, Radium ($Ra$): $0.9$
  • Group 13: Boron ($B$): $2.0$, Aluminum ($Al$): $1.5$, Gallium ($Ga$): $1.6$, Indium ($In$): $1.7$, Thallium ($Tl$): $1.8$
  • Group 14: Carbon ($C$): $2.5$, Silicon ($Si$): $1.8$, Germanium ($Ge$): $1.8$, Tin ($Sn$): $1.8$, Lead ($Pb$): $1.8$
  • Group 15: Nitrogen ($N$): $3.0$, Phosphorus ($P$): $2.1$, Arsenic ($As$): $2.0$, Antimony ($Sb$): $1.9$, Bismuth ($Bi$): $1.9$
  • Group 16: Oxygen ($O$): $3.5$, Sulfur ($S$): $2.5$, Selenium ($Se$): $2.4$, Tellurium ($Te$): $2.1$, Polonium ($Po$): $2.0$
  • Group 17: Fluorine ($F$): $4.0$, Chlorine ($Cl$): $3.0$, Bromine ($Br$): $2.8$, Iodine ($I$): $2.5$, Astatine ($At$): $2.2$
  • Lanthanides series: Range from $1.1$ to $1.3$
  • Actinides series: Range from $1.3$ to $1.5$

Single and Multiple Covalent Bonds

• Bonding Pairs and Lone Pairs:

  • Bonding Pair: A pair of valence electrons shared between two bonded atoms.
  • Lone Pair (Non-Bonding Pair): A pair of valence electrons localized on a single atom that does not participate in chemical bonding.

• Single Bonds:

  • A single covalent bond forms when two atoms share exactly $1$ pair of valence electrons (a total of $2$ electrons).
  • Diatomic Chlorine ($Cl_2$): Each chlorine atom has $7$ valence electrons and shares $1$ electron with the other, forming a single shared pair ($Cl-Cl$) that grants both atoms a full valence shell of $8$ electrons.
  • Hydrogen Chloride ($HCl$): Hydrogen shares its single electron with a chlorine atom sharing one of its seven, yielding a single bond ($H-Cl$) where hydrogen achieves $2$ valence electrons and chlorine achieves $8$.

• Multiple Bonds:
  • Double Bond:
  • Involves the sharing of $2$ electron pairs (a total of $4$ electrons) between two bonded atoms.
  • Carbon Dioxide ($CO_2$): The central carbon atom shares two pairs of electrons with each of the two terminal oxygen atoms, forming two separate double bonds ($O=C=O$).
  • Triple Bond:
  • Involves the sharing of $3$ electron pairs (a total of $6$ electrons) between two bonded atoms.
  • Diatomic Nitrogen ($N_2$): Each nitrogen atom shares $3$ electrons, producing three shared pairs of electrons ($N \equiv N$) and a full octet on each atom.
  • Elemental Restriction for Multiple Bonding:
  • Multiple bonds (double and triple bonds) form almost exclusively among six specific elements: Carbon ($C$), Nitrogen ($N$), Oxygen ($O$), Phosphorus ($P$), Sulfur ($S$), and Silicon ($Si$) — commonly remembered as the $CNOPS\text{ \& }Si$ group.

Covalent Nomenclature and Chemical Formulas

• Numerical Greek Prefixes:

  • Covalent binary compounds use Greek numerical prefixes to denote the number of atoms of each element present:
  • $1$: mono-
  • $2$: di-
  • $3$: tri-
  • $4$: tetra-
  • $5$: penta-
  • $6$: hexa-
  • $7$: hepta-
  • $8$: octa-
  • $9$: nona-
  • $10$: deca-

• Rules for Naming Binary Molecular Compounds:

  1. The first element retains its standard elemental name and receives a Greek prefix corresponding to the number of atoms.
  2. The First-Element Mono- Exception: Never use the prefix mono- on the first element if only one atom is present. All other prefixes (di-, tri-, etc.) must be included on the first element when the subscript is $2$ or greater.
  3. The second element always uses a numerical prefix (including mono- if there is only one atom) and has its suffix modified to -ide.
  4. Vowel Dropping Rule: When combining prefixes with the element oxygen, drop the final "o" from prefixes ending in "o" (mono- becomes mon-) and drop the final "a" from prefixes ending in "a" (tetra- becomes tetr-, penta- becomes pent-, deca- becomes dec-).
  • Example: $CO$ is written as carbon monoxide, not carbon mono-oxide.
  • Example: $N_2O_4$ is written as dinitrogen tetroxide, not dinitrogen tetraoxide.
  • Example: $P_4O_{10}$ is written as tetraphosphorus decoxide, not tetraphosphorus decaoxide.

• Nomenclature Reference Examples:

  • $BCl_3$: Boron trichloride
  • $SF_6$: Sulfur hexafluoride
  • $IF_7$: Iodine heptafluoride
  • $NI_3$: Nitrogen triiodide
  • $N_2O_4$: Dinitrogen tetroxide
  • $Cl_2O$: Dichlorine monoxide
  • $P_4O_{10}$: Tetraphosphorus decoxide
  • $B_5H_9$: Pentaboron nonahydride
  • $Br_3O_8$: Tribromine octoxide
  • $ClF$: Chlorine monofluoride
  • $CO_2$: Carbon dioxide
  • $CO$: Carbon monoxide
  • $S_3O_5$: Trisulfur pentoxide
  • $SiS_2$: Silicon disulfide

• Translating Compound Names to Molecular Formulas:

  • Carbon tetrachloride: $CCl_4$
  • Trisulfur dinitride: $S_3N_2$
  • Silicon dioxide: $SiO_2$
  • Phosphorus triiodide: $PI_3$
  • Diarsenic trioxide: $As_2O_3$

Drawing Lewis Structures for Covalent Compounds

• Step-by-Step Procedure:
  • Step 1: Calculate Total Valence Electrons
  • Sum the valence electrons contributed by each atom present in the molecular formula.
  • Step 2: Determine the Central Atom
  • Apply the following decision hierarchy:
    • Rule a: If carbon ($C$) is present in the formula, carbon is always placed as the central atom.
    • Rule b: If no carbon is present, identify whether there is a unique single atom of an element; that single atom occupies the center.
    • Rule c: If no single atom is present, select the element with the lowest electronegativity value to serve as the central atom.
  • Step 3: Arrange Surrounding Atoms
  • Position all remaining atoms symmetrically around the designated central atom prior to distributing electron pairs and satisfying octets.