Source: Carey, Organic Chemistry, 11th ed., Chapters 1-3
Tags: organic chemistry, structure determines properties, Lewis structures, hybridisation, VSEPR, alkanes, cycloalkanes, conformations, Newman projections, cis-trans isomers, IUPAC nomenclature, UMN CHEM 2301
Difficulty: Foundational | Prerequisites: General Chemistry (CHEM 1062 or equivalent)
Big picture: These three chapters lay the groundwork for everything else in Organic Chemistry I. Chapter 1 establishes how molecular structure (bonding, geometry, electronegativity, intermolecular forces) dictates physical and chemical properties. Chapters 2 and 3 then apply those ideas to the simplest organic molecules, the alkanes and cycloalkanes, covering nomenclature, physical properties, conformational analysis, and geometric (cis-trans) isomerism. If you are joining the course late or feel behind, start here: nearly every later topic assumes fluency with Lewis structures, hybridisation, VSEPR shapes, and Newman projections.
Organic chemistry is built on the idea that a molecule's structure determines how it behaves. Chapter 1 covers the language for describing that structure: Lewis dot structures, formal charge, resonance, orbital hybridisation, molecular geometry, electronegativity, and intermolecular forces. Chapters 2 and 3 introduce the simplest carbon-hydrogen frameworks (alkanes and cycloalkanes), teach you how to name them using IUPAC rules, and then show how rotation around single bonds creates different three-dimensional arrangements (conformations) that differ in energy, and how substituents on rings can be locked into cis or trans relationships.
Organic chemistry
The study of carbon-containing compounds and their properties, reactions, and synthesis. In simple terms, it is the chemistry of carbon-based molecules, which includes nearly all molecules found in living things and most synthetic materials.
Lewis structure (Lewis dot structure)
A diagram showing all valence electrons in a molecule as bonding pairs (lines) and lone pairs (dots), arranged so that each atom achieves a filled valence shell. Think of it as a map of where every electron lives in the molecule.
Formal charge
The charge assigned to an atom in a Lewis structure, calculated as: valence electrons minus (lone-pair electrons + half of bonding electrons). In simple terms, it tells you whether an atom "owns" more or fewer electrons than it brought to the molecule.
Electronegativity
A measure of an atom's ability to attract shared electrons in a covalent bond. Fluorine is the most electronegative element. The practical takeaway: electronegativity differences between bonded atoms create polar bonds.
Polar covalent bond
A bond in which electrons are shared unequally because the bonded atoms differ in electronegativity. Think of it as a tug-of-war where one atom pulls the electron rope harder.
Dipole moment
A quantitative measure of the separation of positive and negative charge in a molecule, reported in debye (D). A molecule can have polar bonds but no net dipole moment if the bond dipoles cancel by symmetry (e.g. CO2).
Resonance structures
Two or more valid Lewis structures for the same molecule that differ only in the placement of electrons (not atoms). The real molecule is a weighted average (the resonance hybrid) of all contributing structures.
Resonance hybrid
The actual electronic structure of a molecule, represented as the blend of all its resonance contributors. No single Lewis structure captures it completely.
Orbital hybridisation
The mixing of atomic orbitals (s, p, and sometimes d) on a single atom to form new, equivalent hybrid orbitals that better describe bonding geometry.
sp3: four hybrid orbitals, tetrahedral geometry, ~109.5 degree bond angles
sp2: three hybrid orbitals, trigonal planar geometry, ~120 degree bond angles
sp: two hybrid orbitals, linear geometry, ~180 degree bond angles
VSEPR (Valence Shell Electron Pair Repulsion)
A model that predicts molecular geometry by assuming that electron groups around a central atom arrange themselves to minimise repulsion. In simple terms, electron clouds push each other as far apart as possible.
Intermolecular forces
Attractive forces between molecules (not within them). Three main types relevant here:
London dispersion forces (present in all molecules; stronger with more surface area and more electrons)
Dipole-dipole interactions (between polar molecules)
Hydrogen bonding (a strong dipole-dipole interaction involving H bonded to N, O, or F)
Functional group
A specific grouping of atoms within a molecule that has predictable chemical behaviour. Think of it as the reactive "handle" on a molecule. Common examples: hydroxyl (-OH), carbonyl (C=O), amino (-NH2).
Constitutional isomers (structural isomers)
Molecules with the same molecular formula but different connectivity of atoms. In simple terms, same atoms, wired up differently.
Carbon forms four bonds because it has four valence electrons and needs four more to complete its octet.
The type of hybridisation determines the geometry around each carbon:
Four single bonds = sp3 = tetrahedral
One double bond + remaining single bonds = sp2 = trigonal planar
One triple bond (or two double bonds) = sp = linear
Sigma bonds (head-on overlap) allow free rotation; pi bonds (side-on overlap of p orbitals) do not.
Every single bond is a sigma bond. A double bond = one sigma + one pi. A triple bond = one sigma + two pi.
Count total valence electrons (adjust for charge).
Connect atoms with single bonds first, then distribute remaining electrons as lone pairs to satisfy octets.
If octets are not satisfied, form double or triple bonds.
Calculate formal charges: atoms should carry the smallest formal charges possible.
Hydrogen is always terminal (it forms only one bond).
Move electrons (lone pairs or pi-bond electrons), never atoms.
All resonance structures must have the same total number of electrons and the same net charge.
The most stable resonance contributor has: the most atoms with complete octets, the fewest formal charges, and any negative formal charge on the most electronegative atom.
Molecules described by resonance are more stable than any single contributor would suggest (resonance stabilisation).
Electronegativity increases going up and to the right in the periodic table (excluding noble gases).
A bond between atoms with different electronegativities is polar. The partial charges are written as delta-plus and delta-minus.
Molecular polarity depends on both bond polarity and molecular geometry. Symmetric molecules can be nonpolar even with polar bonds.
Boiling points and solubility are governed by intermolecular forces:
Nonpolar molecules rely on London dispersion forces; larger surface area = stronger forces = higher boiling point.
Polar molecules also have dipole-dipole interactions, raising the boiling point further.
Hydrogen-bonding molecules (those with O-H, N-H, or F-H) have the highest boiling points for their molecular weight.
"Like dissolves like": polar solvents dissolve polar solutes; nonpolar solvents dissolve nonpolar solutes.
Bronsted acid: proton donor. Bronsted base: proton acceptor.
Lewis acid: electron-pair acceptor. Lewis base: electron-pair donor.
pKa measures acid strength: lower pKa = stronger acid. You will use pKa values throughout the course to predict reaction outcomes.
Factors that stabilise the conjugate base make the parent acid stronger:
Electronegativity of the atom bearing the negative charge
Size of the atom (larger atoms stabilise charge better)
Resonance delocalisation of the negative charge
Inductive effects (electron-withdrawing groups stabilise nearby negative charge)
Hydrocarbon
A molecule composed entirely of carbon and hydrogen. In simple terms, the simplest organic molecules, made only of C and H.
Alkane
A hydrocarbon containing only carbon-carbon single bonds (C-C and C-H). General formula for acyclic alkanes: CnH(2n+2). Also called saturated hydrocarbons because every carbon holds the maximum possible number of hydrogen atoms.
Cycloalkane
An alkane in which the carbon chain forms a ring. General formula: CnH(2n). The ring means two fewer hydrogens compared to the corresponding open-chain alkane.
Homologous series
A family of compounds in which each successive member differs by a CH2 unit. Methane, ethane, propane, butane, and so on form the homologous series of alkanes.
IUPAC nomenclature
The systematic naming system for organic compounds established by the International Union of Pure and Applied Chemistry. This is the naming system you are expected to use in this course.
Constitutional isomers (in context of alkanes)
Alkanes with the same molecular formula but different branching patterns. For example, butane and isobutane (2-methylpropane) are constitutional isomers of C4H10.
Degree of unsaturation (index of hydrogen deficiency)
A calculation that tells you how many rings and/or pi bonds a molecular formula implies. Formula: degrees of unsaturation = (2C + 2 + N - H - X) / 2, where C, N, H, and X are the counts of carbon, nitrogen, hydrogen, and halogen atoms. For alkanes, this value is zero.
Find the longest continuous carbon chain; this gives the parent name (methane, ethane, propane, butane, pentane, hexane, heptane, octane, nonane, decane...).
Number the chain from the end that gives substituents the lowest locants.
Name each substituent (methyl, ethyl, propyl, isopropyl, tert-butyl, etc.) and assign it a number.
List substituents in alphabetical order (ignoring prefixes like di-, tri-, sec-, tert-; but "iso" is part of the name and counts).
Use prefixes di-, tri-, tetra- for identical substituents.
Separate numbers from numbers with commas, numbers from letters with hyphens.
Prefix "cyclo-" before the parent name (cyclopropane, cyclobutane, cyclopentane, cyclohexane...).
If the ring has more carbons than the longest chain substituent, the ring is the parent. If the chain is longer, the ring is a substituent (e.g. cyclohexyl-).
Number the ring to give substituents the lowest set of locants.
Boiling points increase with molecular weight (more electrons = stronger London dispersion forces).
Branching lowers the boiling point because it reduces the surface area available for London dispersion interactions. Pentane (bp 36 C) vs neopentane (bp 9.5 C): same formula, very different boiling points.
Alkanes are nonpolar and insoluble in water. They dissolve in nonpolar solvents.
Density of alkanes is less than water (they float).
Petroleum and natural gas are the primary sources of alkanes.
Combustion of alkanes (reaction with O2 to produce CO2 and H2O) is the basis of fossil-fuel energy.
The combustion equation for any alkane: CnH(2n+2) + (3n+1)/2 O2 -> n CO2 + (n+1) H2O
Conformation
A specific three-dimensional arrangement of atoms in a molecule that results from rotation about a single bond. Different conformations are not different molecules; they interconvert rapidly at room temperature.
Newman projection
A way of viewing a molecule along a particular C-C bond axis. The front carbon is shown as a dot, the rear carbon as a circle. Bonds radiate outward from each.
Staggered conformation
A conformation in which the bonds on adjacent carbons are as far apart as possible (60 degrees apart in a Newman projection). This is lower in energy than eclipsed.
Eclipsed conformation
A conformation in which the bonds on adjacent carbons are directly aligned (0 degrees apart in a Newman projection). This is the highest-energy arrangement due to torsional strain.
Anti conformation
The staggered conformation in which the two largest substituents on adjacent carbons are 180 degrees apart. This is the lowest-energy conformation of butane.
Gauche conformation
The staggered conformation in which the two largest substituents on adjacent carbons are 60 degrees apart. Slightly higher in energy than anti due to steric strain.
Torsional strain
The increase in energy that arises when bonds on adjacent atoms are eclipsed. Thought to result from repulsion between bonding electron pairs.
Steric strain (van der Waals strain)
The increase in energy caused by atoms or groups being forced too close together. Larger groups produce more steric strain.
Angle strain
The strain that arises when bond angles are forced to deviate from their ideal values. Cyclopropane (60 degree angles vs the ideal 109.5 degrees) has severe angle strain.
Ring strain
The combined effect of angle strain and torsional strain in cyclic molecules. Cyclopropane and cyclobutane have significant ring strain; cyclopentane has modest strain; cyclohexane has essentially none.
Chair conformation (of cyclohexane)
The most stable conformation of cyclohexane, in which all bond angles are close to 109.5 degrees and all adjacent C-H bonds are staggered. Cyclohexane adopts this shape to avoid angle and torsional strain entirely.
Axial position
A position on a cyclohexane chair that points straight up or straight down, parallel to the axis of the ring. Three axial bonds point up and three point down, alternating around the ring.
Equatorial position
A position on a cyclohexane chair that points roughly outward from the ring, slightly above or below the equator. Equatorial positions are generally preferred for bulky substituents.
1,3-Diaxial interaction
The steric strain between an axial substituent and the axial hydrogens on the same side of the ring, two carbons away. This is the main reason large substituents prefer the equatorial position.
Ring flip (chair-chair interconversion)
The process by which cyclohexane converts from one chair form to another. Every axial group becomes equatorial, and every equatorial group becomes axial.
Cis-trans isomers (geometric isomers)
Stereoisomers that differ in the arrangement of substituents on the same side (cis) or opposite sides (trans) of a ring or double bond. These are not conformations; they cannot interconvert without breaking bonds.
Rotation about the C-C bond in ethane produces an infinite number of conformations.
The two extremes: staggered (lowest energy) and eclipsed (highest energy).
The energy difference between staggered and eclipsed ethane is approximately 12 kJ/mol (the torsional barrier).
At room temperature, rotation is rapid. Ethane does not "lock" into one conformation.
Looking along the C2-C3 bond, the relevant substituents are the two methyl groups.
Key conformations (from lowest to highest energy):
Anti (methyl groups 180 degrees apart): lowest energy
Gauche (methyl groups 60 degrees apart): about 3.8 kJ/mol higher than anti
Eclipsed with one methyl-methyl eclipse: highest energy (~19 kJ/mol above anti)
The anti and gauche conformations are energy minima (staggered); the eclipsed conformations are energy maxima.
Cyclopropane: very high angle strain (60 vs 109.5 degrees) and torsional strain (all bonds eclipsed). Reactive.
Cyclobutane: high angle strain (90 vs 109.5 degrees). Slight pucker to relieve some torsional strain.
Cyclopentane: mild angle strain. Adopts an "envelope" conformation to reduce torsional strain.
Cyclohexane: no angle strain and no torsional strain in the chair conformation. The most stable common ring.
Drawing a proper chair: start with two parallel, offset horizontal lines, then connect the ends with angled lines to form the chair.
Axial bonds alternate up-down-up-down around the ring.
Equatorial bonds point outward, slightly tilted (each equatorial bond is parallel to the ring bond one position away).
In a ring flip, every axial bond becomes equatorial and vice versa. The "up" or "down" orientation of a substituent does not change during a ring flip.
A single substituent on cyclohexane prefers the equatorial position to avoid 1,3-diaxial interactions.
The larger the substituent, the stronger the preference for equatorial. For tert-butyl, the preference is so strong (~23 kJ/mol) that the ring is essentially locked with tert-butyl equatorial.
For disubstituted cyclohexanes, draw both chairs and assess which places the larger group equatorial, or which minimises total 1,3-diaxial strain.
"Cis" means two substituents are on the same side of the ring plane; "trans" means they are on opposite sides.
Cis and trans isomers of cycloalkanes are different compounds. They cannot interconvert without breaking and reforming bonds.
To assign cis or trans on a chair drawing: "same side" means both up or both down; "opposite sides" means one up and one down.
Cis-1,4 and trans-1,3 disubstituted cyclohexanes can place both substituents equatorial (more stable). Cis-1,3 and trans-1,4 cannot; one group must be axial.
Degree of unsaturation
DoU = (2C + 2 + N - H - X) / 2
Each ring or pi bond contributes one degree. A benzene ring (three pi bonds + one ring) contributes four degrees.
Formal charge
FC = (valence electrons) - (lone pair electrons) - (1/2 bonding electrons)
Combustion of alkanes
CnH(2n+2) + [(3n+1)/2] O2 -> n CO2 + (n+1) H2O
Energy relationships in butane conformations
Anti < gauche (by ~3.8 kJ/mol) < eclipsed-H,CH3 < eclipsed-CH3,CH3 (highest, ~19 kJ/mol above anti)
1,3-Diaxial interaction energies (approximate, per interaction)
Methyl: ~7.6 kJ/mol total for axial vs equatorial
Ethyl: ~8 kJ/mol
Isopropyl: ~9 kJ/mol
tert-Butyl: ~23 kJ/mol (effectively locks the ring)
Conformational analysis is not abstract. The shape preferences of cyclohexane rings govern how steroid hormones (cholesterol, testosterone, oestrogen) behave in your body: every ring junction in a steroid skeleton is either cis or trans, and that geometry controls biological activity.
Alkane combustion is the basis of petrol engines, natural gas heating, and propane grills. The fact that branched alkanes have different physical properties from straight-chain ones is why petrol is a blend of specific hydrocarbons rated by octane number.
Students often think resonance structures represent different molecules that "flip" back and forth. They do not. The molecule is always the hybrid; the individual structures are just our imperfect way of drawing it.
Students frequently confuse constitutional isomers with conformations. Constitutional isomers have different connectivity (different bonds); conformations have the same connectivity but different spatial arrangements from rotation around single bonds.
A common error is numbering the longest chain from the wrong end. Always number so that substituents receive the lowest possible locants.
Many students assume cis is always more stable than trans (or vice versa) on cyclohexane rings. Stability depends on whether the substituents can both be equatorial, which varies with substitution pattern.
⚠️ Drawing correct Lewis structures with proper formal charges is tested heavily. You must be able to assign hybridisation and geometry from a Lewis structure.
⚠️ Resonance is a recurring theme across the entire course. Practice drawing all significant resonance contributors and identifying the most stable one.
⚠️ IUPAC naming questions appear on every exam. Practise naming from structure and drawing from name until it is automatic.
⚠️ Newman projections and conformational energy diagrams for butane are classic exam questions. Be able to draw all staggered and eclipsed forms and rank them by energy.
⚠️ Chair drawings of cyclohexane with axial/equatorial assignments are tested extensively. You must be able to draw both chairs of a substituted cyclohexane, determine which is more stable, and assign cis/trans relationships.
⚠️ Exams are closed-book, short answer and drawing. You may use a model kit piece (one tetrahedral centre, max 5 atoms) to help visualise chirality and conformations.
True or False: A molecule with polar bonds must have a net dipole moment.
Fill in the blank: An sp2-hybridised carbon has ______ geometry and bond angles of approximately ______ degrees.
True or False: The gauche conformation of butane is higher in energy than the anti conformation because of torsional strain.
Fill in the blank: In a cyclohexane ring flip, an axial substituent becomes ______.
True or False: Cis-1,2-dimethylcyclohexane can place both methyl groups in equatorial positions.
Answers: 1. False (bond dipoles can cancel by symmetry). 2. Trigonal planar; 120. 3. False (it is higher due to steric strain, not torsional strain; torsional strain is in eclipsed conformations). 4. Equatorial. 5. False (in cis-1,2, one must be axial and one equatorial).
Q: Draw the most stable Newman projection of 2-methylbutane viewed along the C2-C3 bond. Identify it as anti or gauche.
A: Place the two largest groups (CH3 on C2 and CH2CH3 on C3) anti to each other (180 degrees apart). The remaining positions are occupied by H atoms in staggered arrangements. This is the anti conformation and the most stable rotamer.
Q: Assign the IUPAC name to the following structure: a five-carbon chain with a methyl group on carbon 2 and an ethyl group on carbon 3.
A: 3-Ethyl-2-methylpentane. Substituents are listed alphabetically (ethyl before methyl), and the chain is numbered to give the lowest set of locants.
Q: For trans-1,4-dimethylcyclohexane, draw both chair conformations and identify which is more stable.
A: In trans-1,4, one methyl is "up" and one is "down." In one chair both methyls are axial; in the other both are equatorial. The diequatorial chair is far more stable.
Q: Calculate the degree of unsaturation for C6H12.
A: DoU = (2(6) + 2 - 12) / 2 = (14 - 12) / 2 = 1. One degree of unsaturation, meaning one ring or one double bond.
Q: Explain why neopentane (2,2-dimethylpropane) has a lower boiling point than pentane, despite having the same molecular formula.
A: Neopentane is nearly spherical, minimising the surface area available for London dispersion forces between molecules. Pentane is an elongated chain with more surface contact, producing stronger intermolecular attractions and a higher boiling point.
This material connects directly to Chapter 4 (Chirality), where three-dimensional arrangements become permanent rather than interconvertible, and to Chapter 14 (Spectroscopy), where molecular structure is deduced from spectra. The conformational analysis of cyclohexane is essential for understanding substitution and elimination reactions in Chapters 5-7, because the orientation of a leaving group (axial vs equatorial) affects reaction rates and outcomes.
Lewis dot structure, formal charge calculation, resonance structures, resonance hybrid, orbital hybridisation, sp3 sp2 sp, VSEPR theory, molecular geometry, dipole moment, electronegativity, intermolecular forces, London dispersion, hydrogen bonding, constitutional isomers, structural isomers, degree of unsaturation, index of hydrogen deficiency, alkane, cycloalkane, saturated hydrocarbon, IUPAC nomenclature, systematic naming, Newman projection, staggered, eclipsed, anti, gauche, torsional strain, steric strain, angle strain, ring strain, chair conformation, boat conformation, axial, equatorial, 1-3 diaxial interaction, ring flip, cis-trans isomerism, geometric isomers, conformational analysis, Carey organic chemistry chapter 1, chapter 2, chapter 3, CHEM 2301 UMN, organic chemistry I