Lewis Approach To Chemical Bonding and Covalent Bonds
Kössel-Lewis Approach To Chemical Bonding
The Kössel-Lewis approach, developed independently by Gilbert N. Lewis in the United States and Walther Kössel in Germany around 1916, provided a fundamental conceptual framework for understanding chemical bonding. This approach focused on the role of valence electrons in achieving stable electron configurations.
The central idea is that atoms achieve stability similar to that of noble gases by gaining, losing, or sharing valence electrons.
Octet Rule
The Octet Rule is a key concept introduced by the Kössel-Lewis approach. It states that atoms tend to combine in such a way that they are surrounded by eight valence electrons, a configuration that is electronically stable and similar to that of the noble gases (except Helium, which has a duet).
This stable octet can be achieved through:
- Electron Transfer (Ionic Bonding): An atom loses electrons to achieve the octet of the preceding noble gas, while another atom gains electrons to achieve the octet of the next noble gas. This leads to the formation of ions.
- Electron Sharing (Covalent Bonding): Atoms share valence electrons so that each atom, by sharing, effectively counts eight electrons in its valence shell.
The octet rule is a powerful predictive tool, explaining the formation of many stable chemical compounds. However, it has exceptions, which were later elaborated upon.
Covalent Bond
A covalent bond is a chemical bond formed by the sharing of one or more pairs of valence electrons between two atoms. This sharing allows each participating atom to achieve a more stable electronic configuration, often an octet, in its valence shell.
Formation of Covalent Bonds:
- Covalent bonds typically form between non-metal atoms that have similar electronegativities, meaning neither atom can completely pull electrons away from the other.
- The shared electron pair(s) are attracted by the nuclei of both atoms, holding the atoms together.
Types of Covalent Bonds:
- Single Bond: Formed by sharing one pair of electrons (e.g., H-H in $$H_2$$; Cl-Cl in $$Cl_2$$; C-H in $$CH_4$$). Represented by a single line.
- Double Bond: Formed by sharing two pairs of electrons (e.g., O=O in $$O_2$$; C=C in $$C_2H_4$$). Represented by a double line.
- Triple Bond: Formed by sharing three pairs of electrons (e.g., N≡N in $$N_2$$; C≡C in $$C_2H_2$$). Represented by a triple line.
Polarity of Covalent Bonds:
- If the atoms sharing electrons have different electronegativities, the shared electrons will be pulled more towards the more electronegative atom. This results in a polar covalent bond, where partial positive ($$\delta^{+}$$) and partial negative ($$\delta^{-}$$) charges develop on the atoms.
- If the atoms have identical or very similar electronegativities, the sharing is equal, forming a nonpolar covalent bond (e.g., in $$H_2, O_2, Cl_2$$).
Lewis Representation Of Simple Molecules (The Lewis Structures)
Gilbert N. Lewis introduced a simple notation to represent valence electrons and chemical bonds in molecules. These are called Lewis Structures or Lewis Dot Structures.
Rules for Drawing Lewis Structures:
- Count Total Valence Electrons: Sum the valence electrons of all atoms in the molecule. For polyatomic ions, add electrons for negative charges and subtract for positive charges.
- Determine the Central Atom: The least electronegative atom is usually the central atom (except for Hydrogen, which is always terminal).
- Connect Atoms with Single Bonds: Connect the outer atoms to the central atom with single bonds. Each single bond represents one shared pair of electrons (2 electrons).
- Distribute Remaining Electrons: Distribute the remaining valence electrons as lone pairs around the outer atoms first, to satisfy their octets (or duets for Hydrogen). Then, place any remaining electrons on the central atom.
- Form Multiple Bonds: If the central atom does not have an octet, move lone pairs from surrounding atoms to form multiple bonds (double or triple bonds) between the central atom and the surrounding atoms until the octet rule is satisfied for all atoms (where possible).
- Check Formal Charges (Optional but recommended): Calculate the formal charge on each atom to ensure the most stable Lewis structure is represented. The sum of formal charges should equal the overall charge of the species.
Example: Lewis Structure of Water ($$H_2O$$)
- Total valence electrons: Oxygen (6) + 2 x Hydrogen (1) = 8 valence electrons.
- Central atom: Oxygen (Hydrogen cannot be central).
- Connect atoms: O bonded to two H atoms (uses 2 bonds x 2 electrons/bond = 4 electrons).
- Distribute remaining electrons: 8 - 4 = 4 electrons. Place these 4 electrons as two lone pairs on the Oxygen atom. Each Hydrogen atom has a duet with the single bond.
- Check octets: Oxygen has 2 lone pairs (4 electrons) + 2 bonds (4 electrons) = 8 electrons (octet satisfied). Each Hydrogen has 2 electrons (duet satisfied).
Lewis Structure: $$ \mathrm{H} - \ddot{\mathrm{O}} - \mathrm{H} $$
Formal Charge
Formal Charge is a bookkeeping concept used to determine the distribution of electrons in a Lewis structure. It represents the hypothetical charge an atom would have if all bonds to it were purely covalent and all shared electron pairs were divided equally between the bonded atoms.
Calculation Formula:
$$ \text{Formal Charge} = (\text{Total Valence Electrons of the Atom}) - (\text{Non-bonding Electrons}) - \frac{1}{2}(\text{Bonding Electrons}) $$
Significance of Formal Charges:
- The sum of formal charges on all atoms in a molecule or ion must equal the overall charge of the species.
- A Lewis structure in which the atoms have formal charges closest to zero is generally the most stable and preferred representation.
- When formal charges are non-zero, they are often placed next to the atom to which they belong.
Example: Formal Charges in Ozone ($$O_3$$)
Consider two possible Lewis structures for $$O_3$$, both satisfying the octet rule:
Structure 1: O=O-O
Structure 2: O-O=O
Let's calculate formal charges:
- Structure 1 (O=O-O):
- Left O (double bonded): Valence e⁻ = 6; Non-bonding e⁻ = 4; Bonding e⁻ = 4. FC = 6 - 4 - (4/2) = 6 - 4 - 2 = 0
- Central O: Valence e⁻ = 6; Non-bonding e⁻ = 2; Bonding e⁻ = 6. FC = 6 - 2 - (6/2) = 6 - 2 - 3 = +1
- Right O (single bonded): Valence e⁻ = 6; Non-bonding e⁻ = 6; Bonding e⁻ = 2. FC = 6 - 6 - (2/2) = 6 - 6 - 1 = -1
This structure shows a formal positive charge on the central oxygen and a negative charge on one of the terminal oxygens.
Limitations Of The Octet Rule
While the octet rule is a useful generalization, it does not apply to all molecules and atoms. Several exceptions exist:
1. Incomplete Octets: Some molecules have central atoms with fewer than eight valence electrons.
- Example: Boron trifluoride ($$BF_3$$). Boron has only 6 valence electrons around it in the Lewis structure. $$ \mathrm{F} - \overset{\large \mathrm{F}}{\underset{|}{ \mathrm{B}}} - \mathrm{F} $$
- Example: Beryllium chloride ($$BeCl_2$$). Beryllium has only 4 valence electrons around it.
2. Expanded Octets: Some molecules have central atoms that can accommodate more than eight valence electrons. This is common for elements in the third period and beyond, which have access to d-orbitals.
- Example: Phosphorus pentachloride ($$PCl_5$$). Phosphorus has 10 valence electrons around it. $$ \mathrm{Cl} - \overset{\large \mathrm{Cl}}{\underset{|}{ \mathrm{P}}} - \mathrm{Cl} $$ $$ \qquad \qquad \quad \mathrm{Cl} $$ $$ \qquad \qquad \quad | $$ $$ \qquad \qquad \quad \mathrm{Cl} $$
- Example: Sulfur hexafluoride ($$SF_6$$). Sulfur has 12 valence electrons around it.
3. Molecules with Odd Numbers of Electrons: Some molecules, called free radicals, contain an odd number of electrons, making it impossible for all atoms to satisfy the octet rule.
- Example: Nitric oxide (NO). Nitrogen has 5 valence electrons, Oxygen has 6, and the total is 11. A common Lewis structure shows 7 electrons around N and 8 around O (or vice versa), or some other odd distribution.
4. Elements from Period 3 onwards: Elements like Silicon (Si), Phosphorus (P), Sulfur (S), Chlorine (Cl), etc., can expand their octets because they have access to empty d-orbitals in their valence shell, which can accommodate additional electrons.
5. Certain Compounds of Group 1 and 2 Elements: While metals typically form ionic bonds by losing electrons, some compounds of alkali and alkaline earth metals with highly electronegative elements can have significant covalent character, deviating from simple ionic octet completion.
Despite these exceptions, the Kössel-Lewis approach and the octet rule remain foundational concepts for understanding the basic principles of chemical bonding and the formation of molecules.