Lewis Structures, Bonding Patterns, and Resonance
Lewis Structure: The 4-Step Process
Lewis structures depict the connections between atoms and the distribution of valence electrons in a molecule.
They do not show the three-dimensional shape of a molecule; rotations of a Lewis structure are considered equivalent.
Step 1: Count All Valence Electrons
Sum the valence electrons for all atoms in the molecule.
For ions:
Add electrons for a negative charge (e.g., for \text{N}^{3-}, add 3 electrons).
Subtract electrons for a positive charge (e.g., for \text{NH}_4^+, subtract 1 electron).
Example: For a neutral molecule like \text{PCl}4 (hypothetical, as \text{PCl}4 is actually \text{PCl}4^+, but from the transcript context, assume it's just \text{PCl}4 for electron counting example),
Phosphorus (Group 15) contributes 5 valence electrons.
Each Chlorine (Group 17) contributes 7 valence electrons.
Total valence electrons = 5 + (4 \times 7) = 5 + 28 = 33. (Correcting from transcript's '26' example for PCl3 scenario, where PCl3 is 5 + 37 = 26 electrons). The specific example in transcript for 26 electrons was unclear in the beginning segment, but later it's PCl_x with 26 electrons, implying 4 Chlorine atoms: 5 (P) + 47(Cl) = 33 electrons, so the 26 electron count means the example molecule in transcript has 26 electrons. Let's assume the starting molecule was a PCl3 with an extra lone pair to make 26 electrons for the example, as the molecule PCl4 has 33 for the cation PCl4+. A PCl3 molecule has 26 valence electrons (5 from P + 3*7 from 3 Cl atoms).
Step 2: Connect the Atoms
Identify the central atom(s).
Usually, there is one central atom.
Often, the central atom is listed first in the chemical formula (e.g., P in \text{PCl}_3).
The central atom typically has the lowest electronegativity (though this is a general guideline).
It is possible to have more than one central atom, especially in larger or organic molecules (e.g., \text{H}3 \text{CCH}3 has two central carbon atoms).
Connect the central atom to all outer atoms using single bonds.
Each single bond uses 2 electrons.
Example: For \text{PCl}_3, connect P to each of the three Cl atoms. This uses 3 \times 2 = 6 electrons.
Step 3: Fill Octets
Distribute the remaining electrons as lone pairs to satisfy the octet rule.
Prioritize outer atoms: Place lone pairs on outer atoms until each has an octet (or a duet for hydrogen).
Outer atoms usually cannot expand their octet.
Example: For \text{PCl}_3, after forming 3 single bonds, 26 - 6 = 20 electrons remain. Each Cl needs 6 more electrons to complete its octet (3 lone pairs). So, 3 \times 6 = 18 electrons are placed on the Cl atoms.
Place remaining electrons on the central atom: If any electrons are left after filling outer atom octets, place them on the central atom as lone pairs.
Example: For \text{PCl}_3, after placing 18 electrons on Cl atoms, 20 - 18 = 2 electrons remain. These 2 electrons are placed on the central Phosphorus atom as one lone pair.
Step 4: Check Octets and Exceptions
Verify that all atoms satisfy the octet rule (or its exceptions).
Octet Rule: Most atoms aim for 8 valence electrons (a filled valence shell).
Shared (bonding) electrons are counted twice – once for each atom involved in the bond. (Analogy: money in a bank account is counted by both the bank and the account holder).
Exceptions to the Octet Rule (Less than 8 electrons):
Hydrogen (H): Forms 1 bond, no lone pairs, total 2 electrons (a duet).
Beryllium (Be): Forms 2 bonds, no lone pairs, total 4 electrons.
Boron (B): Forms 3 bonds, no lone pairs, total 6 electrons.
Exceptions to the Octet Rule (More than 8 electrons - Expanded Octets):
Possible for nonmetals from Period 3 onwards (starting with Phosphorus, Sulfur, Chlorine, etc.).
Expanded octets must occur on the central atom.
Addressing Deficiencies: If an atom (not an exception) lacks an octet:
Move a lone pair from an adjacent outer atom (that has an octet) to form a double or triple bond with the central atom.
This increases sharing and helps both atoms achieve octets.
Example for N2 (Nitrogen molecule): Starting with 2 N atoms, 10 valence electrons. Single bond (2 \text{ e}^-, 8 \text{ e}^- remaining). Fill outer Ns (2 \times 6 = 12 \text{ e}^- required, but only 8 \text{ e}^- remain). Each N gets 4 \text{ e}^- (2 lone pairs each), leaving both Ns with 6 electrons. To fix, move a lone pair from each N to form a triple bond, making both Ns have 8 electrons (2\text{lp} + 6_\text{bp} = 8 for each N).
Caution: Ensure that moving lone pairs doesn't cause an outer atom to exceed an octet (unless it's a valid exception, which for outer atoms is rare/non-existent in general rules).
Electronegativity
Electronegativity is an atom's ability to attract electrons in a chemical bond.
Trend: Increases across a period and up a group, generally increasing diagonally towards Fluorine (and Chlorine, as discussed in the context of the diagonal trend).
Common Bonding Patterns (Shortcuts)
Familiarity with common bonding patterns can significantly speed up drawing Lewis structures.
Flowchart Method:
Perform the 'Count' step.
Hypothesize a structure using common bonding patterns for all atoms.
Count the total electrons used in this hypothetical structure.
If the electron count matches the total valence electrons from Step 1, the structure is likely correct.
If it doesn't match, or if any atom cannot adopt its common bonding pattern with the available electrons, revert to the full 4-step process (Connect, Fill, Check).
General Trends: These patterns apply to entire groups of nonmetals on the periodic table (e.g., carbon's pattern applies to silicon, oxygen's pattern applies to sulfur, etc.).
Specific Examples:
Carbon (C) and its group (e.g., Si): Often forms 4 bonds and has 0 lone pairs. These 4 bonds can be: 4 single bonds, 1 triple + 1 single bond, 2 double bonds, or 1 double + 2 single bonds.
Nitrogen (N) and its group: Often forms 3 bonds and has 1 lone pair. These bonds can be: 3 single bonds + 1 lone pair, 1 triple bond + 1 lone pair, or 1 double bond + 1 single bond + 1 lone pair.
Oxygen (O) and its group: Often forms 2 bonds and has 2 lone pairs. These can be: 2 single bonds + 2 lone pairs, or 1 double bond + 2 lone pairs.
Halogens (F, Cl, Br, I): Often forms 1 bond and has 3 lone pairs.
Hydrogen (H): Almost always forms 1 bond and has 0 lone pairs.
Periodic Table Pattern: Moving from carbon across the nonmetals in a period:
Number of lone pairs increases.
Number of bonds decreases.
Resonance Structures
Definition: Resonance occurs when more than one valid Lewis structure can be drawn for a molecule simply by moving electrons (not atoms).
Identification: If you need to fix an octet deficiency by forming a double bond, and there are multiple equivalent positions to place that double bond (e.g., on the left or right side of a central atom), then resonance exists.
Representation: All valid resonance structures are drawn, separated by double-headed arrows (\leftrightarrow).
Example: For \text{SO}_2, one valid structure has a double bond on the left O and a single bond on the right O, while the other has a single bond on the left O and a double bond on the right O. These are shown with a double-headed arrow between them.
The Reality of Resonance:
The molecule does not oscillate between discrete resonance forms. Instead, the true structure is a hybrid or blend of all contributing resonance structures.
Electrons are delocalized over multiple atoms, meaning they are spread out rather than fixed in one location.
This delocalization results in fractional bond orders (e.g., 1.5 bonds for the S-O bonds in \text{SO}_2).
Equivalence: In some cases, all resonance structures may be equivalent (e.g., the two structures for \text{SO}_2), meaning they contribute equally to the overall hybrid structure.
Formal Charge
Purpose: Formal charge is used to evaluate the relative stability and contribution of different Lewis structures, particularly non-equivalent resonance structures, or to compare different isomers.
Counting Method (Different from Octet Rule):
For formal charge, shared (bonding) electrons are counted once – one electron is assigned to each bonding atom.
Each atom