The electrons involved in both ionic and covalent bonds are in the outer levels (valence electrons).

1. Ionic Bonding - electrostatic attraction between ions of opposite charge. Electrons are transferred between atoms.
   
   1. Elements form electron configurations of noble gases. These configurations are very stable, the noble gas configuration has filled an s and a p orbital, thus maintaining the "octet rule." All elements on the periodic table are electrically neutral and have unstable electron configurations (except noble gases, neutral and stable configuration).
      ns²np⁶
      Nonmetals: form anions (negative charge), these elements gain electrons to fulfill the "octet rule" and obtain the next noble gas configuration.
      Metals: form cations (positive charge), lose electrons to form the preceding noble gas.
      Transition Elements: form cations, usually lose the outer s electrons first. Then the transition elements will lose d electrons, following a "general" half-filled orbital rule.
   2. Formulas of Ionic Compounds
      

      Cation	Anions			Examples
      	X-	X-2	X-3
      M+	MX	M2X	M3X	NaF, Na2O, Na3N
      M+2	MX2	MX	M3X2	MgF2, MgO, Mg3N2
      M+3	MX3	M2X3	MX	AlF3, Al2O3, AlN

   3. Sizes of Ions
      

      Na	.186 nm	Cl	.099 nm
      Na+	.095 nm	Cl-	.181 nm

      Cation is smaller than the corresponding metal atom because the excess of protons in the ion draws the outer electrons in closer to the nucleus. In the example above, Na⁺ is smaller than Na.
      Anion is larger than corresponding neutral (chargeless) atom. The extra electron in the anion adds to the repulsion between outer electrons. Also, the added negative charge dilutes the nuclear charge onto more electrons, weakening its effect. In the example above Cl^- is larger than Cl.

   4. Lattice Energy
      The heats of formation of ionic compounds (H). When ionic compounds are formed from the elements, the reaction is always exothermic.
      Lattice energy- the enthalpy change (H), in which the crystal lattice is formed from gaseous ions.
      In other words, it is the measure of the strength of the ionic bond between oppositely charged ions.

2. Covalent Bonding

   Covalent bond - consists of an electron pair shared between two nonmetal atoms.

   -The concept of the covalent bond was first suggested by G.N. Lewis in 1916. Lewis suggested that atoms, by pairing electrons to form an electron-pair bond, could acquire a stable, Noble Gas structure.

   Rules for Lewis Structures
   

   1. Count valence electrons available
      Number of valence electrons contributed by the nonmetal atom is equal to its group number in the periodic table (1 for H). Add electrons to account for negative charge.
      OCl^- ion: 6 (oxygen) + 7 (Chlorine) + 1 (charge) = 14 valence electrons.
   2. Draw a skeleton structure, using single bonds.
      Note that carbon usually forms four bonds with no unshared pairs. In molecules and polyatomic ions, oxygen atoms are ordinarily bonded to a central nonmetal atom rather than to each other. Hydrogen forms only one covalent bond.
      
   3. Deduct two electrons for each single bond in the skeleton.
      OCl^- ion; 14 - 2 = 12 valence electrons left.
   4. Distribute these electrons to give each atom a noble gas structure, if possible.
      If at the end of following steps A-D you find that there are too few electrons to fulfill the octet rule for all atoms makes some electrons so "double duty." Forming a double or triple bond by moving one electron pair corrects a deficiency of electrons. (double bond: two pairs of electrons are shared, triple bond: three pairs of electrons are shared).

      Examples:
      

      OCl-	NO3-	N2

   5. Resonance
      In the NO₃^- ion above, experiment shows that all three bonds are the same length. To explain this, the concept of resonance is used.

      The true structure of NO₃^- is a "hybrid" of these three forms (average of the possible forms)
      Note that:
      

      1. Resonance forms are obtained by moving electrons, not atoms.
      2. Resonance can be expected when it is possible to draw more than one structure that follows the octet rule.
3. Covalent Bond Properties
   Polarity, Bond Energy
   
   1. Polarity Bonds between two identical atoms, as in H-H, Cl-Cl, are nonpolar; the electrons are equally shared between the atoms. If the atoms differ, as in H-F, then the bond is polar; bonding electrons are shifted towards the more electronegative atom. The extent of the polarity depends upon the difference in electronegativity.

      Electronegativity - the ability of an atom in a molecule to attract shared electrons to itself. Based on bond energies, electronegativity is the scale developed by Linus Pauling, ranging from 0.7 (Cs) to 4.00 (F). An electronegativity difference between two atoms of 1.7 results in a bond with 50% ionic character. 3.3 is the maximum difference and is 95% ionic. 0.3 is the minimum (for two different atoms) and is about 10% ionic in character.

      C-H	bond is only very slightly polar (E.N. = .4)
      H-F	bond is strongly polar (E.N. = 1.9)
      Li-F	bond is ionic(E.N. = 3.0)

      
      Electronegativity increases across the periodic table from left to right. It increases going up the periodic table.
      

   2. Bond Energies
      Bond energy values can be used to calculate approximate energies for reactions (H, enthalpy of the reactions). Apply Hess's Law:

      H = (bond energy of bonds broken) - (bond energy of bonds formed)
      H = energy required - energy released.

      The bond energy is defined as H when is broken in the gaseous state.

      Estimation of H from energies. Proceed in two steps;
      

      1. Break up the reactant molecules into atoms. Since bonds are broken
         H₁ = bond energies of the reactants
      2. Combine atoms to form new bonds in products. Since bonds are formed:
         H₂ = bond energies of the products

      In general, multiple bonds are stronger than single bonds.
      Note 1:
      The greater the difference in electronegativity between bonded atoms, the more strongly polar the bond, and the greater the "extra" bond energy.

This page was made by Erik Epp.
Copyright 2002