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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.

    • Viewers are encouraged to subscribe and enable bell notifications to stay updated.

  • 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}2 ext{SO}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.