Chapter 1-8 Organic Chemistry: Lewis Structures, Nomenclature, and Basic Concepts

Course logistics and material access

Gradescope and course materials
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- If there are questions, email the instructor with the problems; otherwise, questions can be addressed in class.
Quick start for today
- The lecture proceeds with how to draw and interpret structures of organic molecules, starting from molecular formulas to Lewis structures and line (skeletal) structures.

Key concepts: Lewis structures vs line structures

Molecular formula basics
- Shows total numbers of carbon (C), hydrogen (H), and heteroatoms in a molecule via subscripts on the symbols.
Lewis structures
- Show valence electrons as dots and bonds as lines; depict all atoms and the electrons involved in bonding.
- Important for visualizing how electrons are distributed and how atoms share electrons.
Line (skeletal) structures
- Simplified representation; bonds are shown as lines and the carbon skeleton is implied by vertices.
- Hydrogens attached to carbons are generally not drawn explicitly in line structures; heteroatoms and any hydrogens bound to them are shown.
Relationship between the three representations
- Molecular formula provides counts (C, H, heteroatoms).
- Lewis structure explicitly shows valence electrons and bonds.
- Line structure shows connectivity with fewer details about every electron.

Valence electrons, octet rule, and periodic table refresher

Valence electrons
- The electrons involved in bonding; for Lewis structures you must know the valence electrons for each atom in the molecule.
Octet rule
- Most elements (except hydrogen) prefer to have eight electrons in their outer shell (octet) for stability.
- Hydrogen is an exception: it needs 2 electrons to be stable.
- Helium already has a full outer shell with 2 electrons (no further discussion needed here).
Quick periodic table reference (in the context of Lewis structures)
- The s-block and p-block determine valence electron counts; noble gases (Group 18) are generally nonreactive.
- Halogens (Group 17) have seven valence electrons and tend to gain or share one electron to complete their octet.
- Main takeaway: to build Lewis structures, know the valence electrons for the atoms involved and use octet completion as a guide.
Visualizing valence electrons by group (simplified)
- Hydrogen: 1 valence electron
- Group 2 (e.g., Be): 2 valence electrons
- Group 13 (3A): 3 valence electrons
- Group 14 (4A): 4 valence electrons
- Halogens (Group 17): 7 valence electrons
- Noble gases (Group 18): full octet (except He, which has 2)

Example: methane (CH₄)

Valence electron counts
- Carbon: 4 valence electrons
- Hydrogen: 1 valence electron each (×4)
- Total valence electrons: 4 + 4 imes 1 = 8
Central atom choice
- Carbon is the central atom because it can form the maximum number of bonds (4).
Electron distribution to satisfy octets
- Carbon needs 4 more electrons; each hydrogen needs 1 more electron.
- Place carbon at center; connect four hydrogens with single bonds (C–H).
- Each H has 2 electrons in its bond (1 from itself, 1 from carbon); carbon shares electrons to achieve 8 around it.
Bond representation
- Bonds between carbon and hydrogen are shown as lines after electrons are paired.
Formula and classification
- This saturated molecule follows the alkane general formula: CnH{2n+2} with n = 1 for methane, giving CH_4.
Summary of the methane Lewis structure
- Carbon in the center with four C–H single bonds; each atom satisfies its octet (H has 2 electrons in its shell, C has 8 around it).

Ethane (C₂H₆) and unsaturation concept

Building ethane (two carbons)
- Central atoms: two carbons bonded to each other; each carbon also bonded to hydrogens to satisfy valence.
- Each carbon typically forms three C–H bonds in CH₃ groups plus one C–C bond to the other carbon.
- Resulting formula: ext{CH}3- ext{CH}3; overall: ext{C}2 ext{H}6.
Saturated vs unsaturated
- Saturated molecules have only single bonds between carbon atoms (alkanes).
- If there is a C=C double bond or a C≡C triple bond, the molecule is unsaturated.
- Example of unsaturation: a double bond between two carbons reduces the number of hydrogens by 2 compared to the saturated alkane with the same number of carbons.

Carbonyl compounds: ketones, aldehydes, and carboxylates

Carbonyl group basics
- Carbonyl group: conjugated carbon-oxygen double bond: ext{C}= ext{O}
- Compounds containing C=O can be categorized by where the carbonyl carbon is connected.
Ketones vs aldehydes
- Ketone: carbonyl carbon is bonded to two other carbons (R–CO–R′).
- Example: acetone, ext{CH}3- ext{CO}- ext{CH}3 (a ketone).
- Aldehyde: carbonyl carbon is bonded to one carbon and one hydrogen (R–CHO).
- Example: acetaldehyde, ext{CH}_3- ext{CHO}.
Carboxylate and carboxylic acids (brief overview mentioned in the lecture)
- Carboxylic acids: R–COOH (carbonyl carbon bonded to an OH group).
- Deprotonated form (carboxylate) can carry a negative charge and show resonance; common example: acetate anion in basic conditions.
- Ionization in water and reactivity of carboxylic acids/acids bases will be covered later.

Line structures: rules for drawing and interpretation

Why line structures are common
- Easier to draw for large molecules; still conveys connectivity.
Rules for carbon and hydrogen in line structures
- Vertices (where lines meet) represent carbon atoms; hydrogens attached to carbons are usually not drawn explicitly.
- When a heteroatom is present, it is shown in the structure along with the carbon it's bonded to, and hydrogens attached to the heteroatom are shown (e.g., OH groups).
Heteroatoms and lone pairs
- Heteroatoms (anything other than C and H in organic compounds: O, N, S, P) are shown; their lone pairs may be indicated in Lewis structures.
Hydrogen on heteroatoms
- OH group is shown as an explicit bond to H; the hydrogen on heteroatoms is usually shown in line structures.
Special cases with charges
- If a carbon lacks electrons or has an extra electron, charges may appear as + or − signs.
- Example: carbanion (negative charge) vs carbocation (positive charge, sometimes historically called carbonium ion).
Example kinds of charged situations
- If a carbon has an extra electron pair, it may carry a negative charge (carbanion).
- If a carbon is short of electrons (three bonds instead of four) and carries a positive charge, it can be a carbocation.

Naming and substituents: hydrocarbons and basic IUPAC concepts

Straight-chain alkanes (saturated hydrocarbons)
- Methane, ethane, propane, butane, pentane, hexane, heptane, octane, nonane, decane, etc.
- General formula: CnH{2n+2} for alkanes.
- Common names (for quick recall): methane (CH₄), ethane (C₂H₆), propane (C₃H₈), butane (C₄H₁₀), etc.
Substituent naming (alkyl groups)
- When a carbon chain has a side group, that side group is named as a substituent: methyl (CH₃–), ethyl (C₂H₅–), propyl (C₃H₇–), etc.
- Examples: a methyl substituent on a chain is called a methyl group; an ethyl substituent is an ethyl group.
Ring (cyclo) compounds
- Cyclopropane, cyclobutane, cyclopentane, cyclohexane, cyclooctane, etc.
- Prefix cyclo- indicates a ring structure; these are cycloalkanes and share the same saturated hydrocarbon framework but in a ring.
Carbon classification by bonding environment (primary/secondary/tertiary/quaternary)
- Primary carbon: bonded to one other carbon (CH₃– group at a terminus is a primary carbon).
- Secondary carbon: bonded to two other carbons.
- Tertiary carbon: bonded to three other carbons.
- Quaternary carbon: bonded to four other carbons (no hydrogens).
- These classifications help in understanding reactivity and naming conventions.
Practical notes on naming in practice
- When a carbon skeleton is named, substituents (methyl, ethyl, etc.) are named as prefixes to the main chain or the ring, with the suffix -ane for alkanes.
- The instructor emphasizes practicing the basic set up to avoid common naming mistakes, and that many molecules can be drawn from a given formula by choosing possible connectivities; however, the core rules (saturation, ring formation, substituents) guide correct enumeration.

Quick practical reminders and synthesis of ideas

Structural representation choices
- Lewis structures provide the most complete picture of electrons and bonding.
- Line structures are typically used for readability in larger molecules, with implicit hydrogens on carbons and explicit heteroatoms when present.
Octet rule and exceptions
- Most elements strive for an octet, but hydrogen follows the duet rule (2 electrons).
When to expect charges in structures
- If valence electrons do not add up to satisfy a neutral valence state, charges may appear on carbons (carbocation or carbanion) or other atoms as needed.
Connecting theory to real molecules
- Methane CH₄ shows the basic saturation concept and the C–H single bonds.
- Ethane adds a second carbon and a C–C single bond.
- Ketones and aldehydes introduce carbonyl chemistry (C=O) and the difference between ketones (R–CO–R′) and aldehydes (R–CHO).
- Alcohols and other heteroatom-containing molecules are treated by showing heteroatoms and their hydrogen attachments clearly in line structures.
Studying strategy suggested by the instructor
- Build intuition by starting from the molecular formula, determine the central atom(s), then distribute electrons to satisfy octets, and represent bonds accordingly.
- Practice naming with methane to decane for straight chains, and then add cycloalkanes and substituents to become proficient in IUPAC naming conventions.

Illustrative equations and formulas (recap)

Alkane general formula
- CnH{2n+2}
Carbonyl basics (illustrative)
- Ketone: R-CO-R' with a carbonyl group C=O; example acetone: ext{CH}3- ext{CO}- ext{CH}3
- Aldehyde: R-CHO; example acetaldehyde: ext{CH}_3- ext{CHO}
Carboxylic acids and related ions (brief)
- Carboxylic acid: R-COOH
- Carboxylate anion (example): R-COO^-
General representation of carbon and hydrogen in skeletal structures
- Methane: ext{CH}_4
- Ethane: ext{C}2 ext{H}6
- Propane: ext{C}3 ext{H}8
- Cycloalkanes: cyclopropane, cyclobutane, cyclopentane, cyclohexane, cyclooctane

Quick reference cheat sheet (summary)

Valence counts to memorize (quick guide)
- Hydrogen: 1 valence electron
- Carbon: 4 valence electrons
- Halogens: 7 valence electrons
- Noble gases: full valence shells (8 or full outer shell, He with 2)
Saturation rule
- Single bonds only between carbons: saturated hydrocarbon (alkane)
- Double/triple bonds indicate unsaturation (alkenes, alkynes)
Naming anchors
- Methane, Ethane, Propane, Butane, Pentane, Hexane, Heptane, Octane, Nonane, Decane
- Methyl, Ethyl as substituents; cyclo- prefix for rings
Structural conventions to apply in exam-style problems
- Use Lewis structures for electron distribution; switch to line structures for large molecules; remember that hydrogens on carbons are implicit in line drawings, while heteroatoms and their hydrogens are explicit.
- Check for charges if octet is not satisfied; identify carbocation or carbanion as needed.