Chemical Bonds and Electronegativity

• Atoms acquire complete valence shells in 2 ways: ionic bonds and covalent bonds

• Electronegativity: a measure of an atom’s attraction for electrons that it shares in a chemical bond with another atom

  • Increases left to right within a row/period

  • Decreases from top to bottom within a column

• First ionization potential increases from left to right, and decreases from top to bottom because the principle quantum levels increase and the outermost electrons are farther from the nuclei (held less tightly)

• Electron affinity: energy released upon addition of an electron

  • Becomes more favorable as you go left to right (closer to the noble gas configuration)

  • Decreases from top to bottom

Formation of Ions

• Ions are formed by the transfer of electrons from the valence shell of an atom of electronegativity to the valence shell of an atom of higher electronegativity

  • Ions form if the difference in EN is 1.9 or greater

  • Results in ions with the same electron configuration as the noble gas nearest to there atomic number

  • The attraction between ions allows ionic salts to form a strong crystal lattice and high melting point

Covalent Bonds

• Covalent bonds are formed between atoms that share one or more pairs of electrons to give a noble gas configuration to each atom

• Lewis model: electrons in covalent bond functions in 2 ways simultaneously → shared between 2 atoms and fills the outer valence of each

• Amount of energy released when the bond forms is the same amount required to break the bond

• System becomes more stable when covalent bonds are formed because the wave character of the electrons is stabilized relative to 2 separate atoms

• Bond length: the distance between nuclei participating in a covalent bond

  • Every covalent bond has a characteristic bond length

• Each bond requires 2 electrons

• For every lone pair on an atom, one fewer bond is possible

Polar Covalent Bonds

• Polar covalent bonds: share electrons unequally

  • Polarity in the bond increases with increasing difference in electronegativity between the bonded atoms

  • The more EN atom gains a greater fraction of the shared electrons and acquires a partial negative charge

  • The less EN atoms has a smaller fraction of the shared electrons and acquires a partial positive charge

    • Both partial charges are the same in absolute magnitude

  • The direction of bond polarity is shown with an arrow, the arrowhead pointing toward the neg end and a plus sign by the pos end

• Less than 0.5 = 0.5 nonpolar covalent

• 0.5 to 1.9 = 1.9 polar covalent

• Greater than 1.9 = 1.9 ions formed

• EN varies somewhat depending on the chemical environment and oxidation state of  an atom

• Bond dipole moment: a vector quantity that measures the polarity of a covalent bond

  • The charge on either atom in the bond x the distance between the nuclei

    • e x d

Guidelines for Writing Lewis Structures

• Determine the number of valence electrons in the molecule or ion

  • Add the number of valence electrons contributed by each atom (for ions add one electron for each neg charge and subtract one electron for each pos charge)

• Determine the arrangement of atoms

  • Except for the simplest molecules and ions, this is determined experimentally because alternative possibilities may lead to isomers

    • Isomers: same molecular formula, diff arrangement of atoms

• Connect the atoms with single bonds; then arrange the remaining electrons in the pair so that each atom in the molecule or ion has a complete outer shell

• Each pair of electrons shared between two atoms is shown as a single line between the atoms; each unshared pair of electrons is shown as a pair of dots

• If 2 atoms share only a single pair of electrons, they form a single bond; 2 pairs = 2 bonds, etc

• Bonding electrons: valence electrons involved in forming a covalent bond

• Nonbonding electrons: valence electrons not involved in forming covalent bonds

• Formal charge: the charge on an atom in a polyatomic ion or molecule

• To determine formal charge

  • Write a correct lewis structure for the molecule or ion

  • Formal charge = # of valence electrons in the neutral nonbonded atom - (unshared electrons + 1/2shared electrons)

  • The sum of all formal charges is equal to the total charge on the molecule or ion

  • Any atoms with an octet plus one more bond than its neutral form has a positive charge