Organic Chemistry Chapter 1: Lewis Structures

Chapter 1: Review of General Chemistry

• This first chapter of the organic chemistry playlist will cover several fundamental topics as a review from general chemistry, including:

• Lewis structures (the focus of this lesson)

• Formal charges

• Hybridization and valence bond theory

• Molecular orbital theory (acknowledged as often 'skimped on' in general chemistry, potentially requiring more focused review)

• Polarity and intermolecular forces

Introduction to the Organic Chemistry Playlist

• The instructor will be using a whiteboard for instruction, a direct response to student preference over PowerPoint and green screen.

• New lessons will be released weekly throughout the $2020-2021$ school year.

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Introduction to Organic Chemistry

• Organic chemistry is defined as the study of molecules found in living systems, contrasting with inorganic chemistry.

• Historically, it was believed that a "vital force" was involved, but this is not relevant to the modern definition.

• The core focus of organic chemistry is the study of carbon-containing compounds.

• Unlike inorganic chemistry, which encompasses all other elements, organic chemistry primarily deals with carbon and its interactions.

• This focus simplifies the study, as it largely excludes transition metals.

Variability in Organic Chemistry Instruction

• Organic chemistry is not taught uniformly across textbooks, professors, or universities.

• The instructor aims to provide a "middle-of-the-road" course, informing students when there is variability in presentation or rules.

• Real-world chemical reality is complex; simplified trends are often taught as absolute rules without full disclosure.

• The instructor will explicitly highlight areas where different presentations exist and advise students to note their professor's approach.

• Such instances will not be super common but will occur in key places.

Lewis Structures: Valence Electrons

• Lewis structures represent valence electrons, which are determined by an element's group on the periodic table:

• Group $1$ (e.g., hydrogen): $1$ valence electron.

• Group $2$ (e.g., beryllium): $2$ valence electrons.

• Boron's column: $3$ valence electrons.

• Carbon's column: $4$ valence electrons.

• Nitrogen's column: $5$ valence electrons.

• Oxygen's column: $6$ valence electrons.

• Halogens: $7$ valence electrons.

• Noble gases: $8$ valence electrons (generally not involved in bonding relevant to this chapter).

Predicting Typical Number of Bonds

• The typical number of bonds an atom forms can be predicted by how many electrons it needs to achieve a filled octet.

• Halogens ($7$ valence electrons) need $1$ more, so they typically make $1$ bond.

• Oxygen's column ($6$ valence electrons) needs $2$ more, so they typically make $2$ bonds.

• Nitrogen's column ($5$ valence electrons) needs $3$ more, so they typically make $3$ bonds.

• Carbon's column ($4$ valence electrons) needs $4$ more, so they typically make $4$ bonds.

• Boron's column ($3$ valence electrons) typically makes $3$ bonds, as making $5$ would violate the octet rule.

The Octet Rule

• Most atoms strive to have $8$ valence electrons to achieve stability.

• Origin of the Octet Rule: A typical electron shell contains an $s$ subshell (one orbital, $2$ electrons) and a $p$ subshell (three orbitals, $6$ electrons), totaling $8$ positions for electrons ($2 + 6 = 8$).

Exceptions to the Octet Rule

• Expanded Octet (More than 8 electrons):

• Occurs in elements from the third period and lower (e.g., sulfur), not in elements like carbon, nitrogen, or oxygen, which are the main focus of organic chemistry.

• Reason: Starting in the third shell, there is also a $d$ subshell, which can accommodate up to $10$ additional electrons.

• Example: In sulfuric acid ($ext{H}<em>2 ext{SO}</em>4$), sulfur can be shown with $12$ electrons around it ($2, 4, 6, 8, 10, 12$).

• While third-period elements can go over the octet, they don't have to.

• Under the Octet Rule (Less than 8 electrons):

• Hydrogen: Only wants $2$ electrons (a duet) to resemble helium, as its first shell only has an $s$ subshell.

• Beryllium: (more metallic, less common in covalent bonding but possible in review) Typically forms $2$ bonds, resulting in only $4$ electrons around it (e.g., in $ext{BeH}_2$).

• Boron and Aluminum: Typically form $3$ bonds, resulting in only $6$ electrons around them (e.g., in $ext{BH}_3$). It is considered normal for them to have an incomplete octet.

• Odd Number of Valence Electrons: (Highly unlikely in organic chemistry)

• Species with an overall odd number of valence electrons (e.g., $ext{NO}$) will necessarily have at least one atom with an odd number of electrons, thus violating the octet rule.

Methods to Achieve a Filled Octet

• Ionic Bonding (Metal + Non-metal):

• Involves the complete transfer of electrons.

• Example: $ext{Na}$ (1 valence electron) and $ext{Cl}$ (7 valence electrons).

  • $ext{Na}$ donates its electron to $ext{Cl}$. Result: $ext{Na}^+$ and $ext{Cl}^-$.

  • $ext{Cl}^-$ now has $8$ valence electrons. $ext{Na}^+$ effectively has a filled octet from its complete inner shell.

• Less common in the study of organic chemistry.

• Covalent Bonding (Two Non-metals):

• Involves the sharing of electrons.

• Example: Two $ext{Cl}$ atoms.

  • Neither $ext{Cl}$ easily loses an electron; both want to gain.

  • They share a pair of electrons, forming a covalent bond (represented as a line, $ext{Cl}- ext{Cl}$).

  • Both $ext{Cl}$ atoms count the shared pair towards their octet, achieving $8$ electrons each.

• This is the predominant type of bonding encountered in organic chemistry.

General Steps for Drawing Lewis Structures

1. Count Total Valence Electrons: Sum the valence electrons from all atoms in the molecule.

2. Identify Central Atom(s): Usually the atom that can make the most bonds (not hydrogen). Often the least electronegative atom (excluding hydrogen).

3. Set Up the Skeleton: Connect the central atom(s) to the surrounding atoms using single bonds.

4. Fill Outside Atoms' Octets: Distribute remaining electrons as lone pairs to the outer atoms first (except hydrogen, which only needs a duet or $2$ electrons).

5. Place Remaining Electrons on Central Atom: Any electrons left after filling outside atoms go onto the central atom(s) as lone pairs.

6. Check Central Atom Octet and Form Multiple Bonds if Necessary: If the central atom lacks a filled octet and there are no more electrons to add, convert lone pairs from adjacent atoms into multiple bonds with the central atom until its octet is satisfied.

Lewis Structure Examples

• Methane ($ext{CH}_4$):

• Total valence electrons: C ($4$) + $4$ H ($4 imes 1 = 4$) = $8$.

• Central atom: Carbon.

• Skeleton: Carbon bonded to $4$ hydrogens (single bonds).

• All $8$ electrons used (C-H bonds). All hydrogens have $2$, carbon has $8$. Octets satisfied.

• Ammonia ($ext{NH}_3$):

• Total valence electrons: N ($5$) + $3$ H ($3 imes 1 = 3$) = $8$.

• Central atom: Nitrogen.

• Skeleton: Nitrogen bonded to $3$ hydrogens (single bonds), uses $6$ electrons.

• $2$ electrons remaining. Place as a lone pair on nitrogen.

• All $8$ electrons used. Hydrogens have $2$, nitrogen has $8$ (six bonding, two non-bonding). Octets satisfied.

• Formaldehyde ($ext{H}_2 ext{CO}$):

• Total valence electrons: $2$ H ($2 imes 1 = 2$) + C ($4$) + O ($6$) = $12$.

• Central atom: Carbon.

• Skeleton: Carbon bonded to $2$ hydrogens and $1$ oxygen (single bonds), uses $6$ electrons.

• $6$ electrons remaining. Outside atoms: Hydrogens are full. Oxygen needs $6$ (put $3$ lone pairs on oxygen). All $12$ electrons used.

• Carbon has only $6$ electrons. Oxygen has a lone pair to share.

• Convert one lone pair from oxygen to form a double bond with carbon.

• Now carbon has $8$ electrons ($2$ from H, $2$ from H, $4$ from O=C). All octets satisfied.

• Acetaldehyde ($ext{CH}_3 ext{CHO}$) (Example of a condensed formula):

• Total valence electrons: $2$ C ($2 imes 4 = 8$) + O ($6$) + $4$ H ($4 imes 1 = 4$) = $18$.

• From condensed formula, first carbon is bonded to $3$ H, then to second carbon. Second carbon is bonded to $1$ H and $1$ O.

• Skeleton: $ext{C}( ext{H}_3)- ext{C}( ext{H})- ext{O}$ (single bonds initially). Uses $2(3) + 2(1) + 2(1) + 2(1) = 12$ electrons.

• $6$ electrons remaining. Hydrogens are full. Put $3$ lone pairs on oxygen. All $18$ electrons used.

• First carbon has $8$ electrons. Second carbon has only $6$. Oxygen has lone pairs.

• Convert one lone pair from oxygen to form a double bond with the second carbon.

• Now all carbons and oxygen have $8$ electrons. All atoms are making their typical number of bonds (C: $4$, O: $2$, H: $1$).

• Acetic Acid ($ext{CH}_3 ext{COOH}$) (Carboxylic acid, common organic molecule):

• Total valence electrons: $2$ C ($2 imes 4 = 8$) + $4$ H ($4 imes 1 = 4$) + $2$ O ($2 imes 6 = 12$) = $24$.

• Initial, incorrect skeletal attempt (e.g., $ext{C}( ext{H}_3)- ext{C}- ext{O}- ext{O}- ext{H}$) is discussed and rejected because it would lead to atoms not making their typical number of bonds. This often indicates formal charges and instability (a topic for the next lesson).

• Correct Skeleton: The second carbon is bonded to both oxygens, and one of those oxygens is bonded to the hydrogen. $ext{C}( ext{H}_3)- ext{C}- ext{O}( ext{H})$ and another $ext{O}$ also bonded to that second carbon.

• Skeleton: First carbon bonded to $3$ hydrogens, then to second carbon. Second carbon bonded to both oxygens. One oxygen bonded to hydrogen. Uses $2(3) + 2(1) + 2(1) + 2(1) = 16$ electrons.

• $8$ electrons remaining. Fill outside atoms (hydrogens are full, give $3$ lone pairs to each oxygen). All $24$ electrons used.

• First carbon has $8$ electrons. Second carbon has only $6$. Both oxygens have lone pairs. The single-bonded oxygen is only making $1$ bond (vs. typical $2$).

• Convert one lone pair from the single-bonded oxygen to form a double bond with the second carbon.

• Now all atoms are making their typical number of bonds (first C: $4$, second C: $4$, double-bonded O: $2$, single-bonded O: $2$, H: $1$). This is the proper and more stable Lewis structure.

• Key Principle: If an atom is not making its typical number of bonds, it suggests instability and possible formal charges. Structures where all atoms maintain their typical bonding patterns are generally preferred.

• Recognizing condensed formulas like $ext{COOH}$ will require memorization of their associated Lewis structures by chapter two.

Conclusion

• This lesson covered Lewis structures, mostly as a review from general chemistry.

• Additional practice and study guides are available on the instructor's premium course website.