Difficulty: Intermediate | Prerequisites: General Chemistry (bonding, electronegativity, Lewis structures), Chapter 7 (functional groups basics)
Chapter 8 is where organic chemistry stops being purely about structure and starts asking: how do molecules behave in the physical world? You are learning how the type of carbon a halogen is bonded to (sp³, sp², aromatic) classifies the halide, and then how the intermolecular forces between molecules govern measurable properties like boiling point, melting point, and solubility. If you have a solid grip on electronegativity, Lewis structures, and the basics of molecular shape from general chemistry, you are ready for this. The classification of halides here feeds directly into the substitution and elimination reactions in Chapters 9 and 10, so getting comfortable with this vocabulary now saves real pain later.
Alkyl, vinylic, and aryl halides differ by the hybridisation of the carbon bearing the halogen. Intermolecular forces (London dispersion, dipole-dipole, and hydrogen bonding) determine boiling points, melting points, and solubility. "Like dissolves like": molecules need similar intermolecular forces to mix favourably, and water solubility requires the ability to both donate and accept hydrogen bonds.
Alkyl halide
A compound in which a halogen (F, Cl, Br, or I) is directly attached to an sp³-hybridised carbon of an alkyl group (C–C single-bond framework). Think of it as the simplest case: a halogen hanging off a fully saturated carbon.
Vinylic halide
A halide where the halogen is bonded to an sp²-hybridised carbon that is part of a non-aromatic C=C double bond. In simple terms, the halogen sits right on the double bond, not next to it.
Aryl halide
A halide where the halogen is directly bonded to an sp²-hybridised carbon of a benzene ring or other aromatic ring. Think of it as a halogen plugged into an aromatic system.
Intermolecular forces (IMFs)
Forces of attraction between separate molecules (as opposed to intramolecular bonds within a molecule). These govern physical properties like boiling point, melting point, and solubility.
London dispersion forces (LDF)
Weak, temporary attractive forces that arise from instantaneous dipoles caused by electron movement. Present in all molecules. In simple terms, electrons slosh around and create fleeting charge imbalances that neighbouring molecules respond to.
Polarisability
A measure of how easily a molecule's electron cloud can be distorted. Proportional to the number of electrons and greater in larger, more linear molecules with more surface area. Think of it as how "squishy" the electron cloud is.
Dipole-dipole interactions
Attractive forces between the partial positive end of one polar molecule and the partial negative end of another. Strength depends on the magnitude of the partial charges and anti-symmetry (asymmetry) of the molecule.
Hydrogen bonds
A particularly strong type of dipole-dipole interaction that occurs when H is bonded to a highly electronegative atom (N, O, or F) and interacts with a lone pair on another N, O, or F. Typically the strongest IMF in organic molecules.
Miscible
Two liquids that mix in all proportions, forming a single phase. Think of it as "always mixes, no matter the ratio."
ΔG (Gibbs free energy of mixing)
ΔG = ΔH – TΔS. If ΔG < 0, mixing is thermodynamically favourable. If ΔG > 0, mixing is unfavourable. ΔS of mixing is always positive (mixing increases disorder), so the sign of ΔG depends heavily on ΔH.
"Like dissolves like"
The principle that substances with similar intermolecular forces tend to be soluble in one another. Polar dissolves polar; nonpolar dissolves nonpolar.
The classification turns entirely on the hybridisation of the carbon that holds the halogen.
Alkyl halide: Halogen on an sp³ carbon, part of a C–C single-bond chain. These are the ones you will see most in SN1, SN2, E1, and E2 reactions.
Vinylic halide: Halogen on an sp² carbon of a C=C double bond (not aromatic). These are far less reactive toward nucleophilic substitution because the C–X bond has partial double-bond character and the carbon is harder to access.
Aryl halide: Halogen on an sp² carbon of a benzene or other aromatic ring. Similar to vinylic halides in resistance to standard substitution, but they react under specialised conditions (e.g. SNAr, benzyne mechanism).
The key take-away: sp³ carbons holding halogens behave very differently from sp² carbons holding halogens. Most of Chapters 9 and 10 focus on reactions of alkyl halides specifically.
Three types matter here, in order of typical increasing strength:
London dispersion forces (LDF)
Present in every molecule, polar or not.
Arise from temporary, instantaneous dipoles as electrons move.
Strength scales with polarisability, which is proportional to the number of electrons (roughly molecular weight).
Greater in linear molecules than branched ones of the same mass, because a linear shape exposes more surface area for contact.
Dispersion forces are the dominant IMF for large, nonpolar molecules.
Dipole-dipole interactions
Occur between polar molecules that have a permanent uneven distribution of charge.
Strength depends on the size of the partial charges and how asymmetric the molecule is (anti-symmetry).
A symmetric molecule can have polar bonds but no net dipole (e.g. CF₄), so dipole-dipole interactions may be absent even when polar bonds are present.
Hydrogen bonds
A special, stronger case of dipole-dipole interaction.
Requires H bonded to N, O, or F (the H–X dipole strength matters).
The hydrogen interacts with a lone pair on N, O, or F of a neighbouring molecule.
Typically the strongest IMF you will encounter in organic chemistry.
Oxygen is more electronegative than nitrogen, so O–H···O hydrogen bonds are generally stronger than N–H···N ones. A small difference in electronegativity between N and O translates to a big difference in hydrogen-bond strength.
The ranking of these forces determines which molecule has the higher boiling point, the higher melting point, and the greater solubility in a given solvent.
Boiling point and enthalpy/entropy
Boiling requires overcoming intermolecular attractions (ΔH > 0) while gaining freedom of movement (ΔS > 0).
At the boiling point, ΔG = 0, so T(bp) = ΔH(vap) / ΔS(vap).
Because ΔS of vaporisation is roughly constant across similar molecules, boiling point is approximately proportional to ΔH(vap). Stronger IMFs mean higher ΔH(vap), which means a higher boiling point.
Ranking boiling points: the decision process
Compare hydrogen bonding first. A molecule that can hydrogen-bond generally boils higher than one that cannot, all else being similar.
If hydrogen bonding is the same (or absent in both), compare dipole-dipole. Larger net dipole means a higher boiling point.
If polarity is similar, compare London dispersion. More electrons (higher molecular weight) and more surface area (linear > branched) mean stronger LDF and higher boiling point.
Worked example from the notes: CH₃CH₃ vs CH₃CF₃ vs CF₃CF₃
CH₃CH₃: no dipole, smallest molecular weight. Lowest boiling point.
CF₃CF₃: heavy, but highly symmetric with extremely low polarisability. Despite high molecular weight, boiling point is lower than expected.
CH₃CF₃: moderate weight, has a significant dipole (asymmetric). Highest boiling point of the three.
This example illustrates that molecular weight and electron count are the most important metrics for dispersion forces, but polarity and polarisability can override a simple mass-based ranking.
\Delta G = \Delta H - T \Delta SAt the boiling point, ΔG = 0:
T_{bp} = \frac{\Delta H_{vap}}{\Delta S_{vap}}Because ΔS(vap) is roughly constant, boiling point tracks ΔH(vap) almost directly.
For solubility (Gibbs free energy of mixing):
\Delta G_s = G(\text{solution}) - [G(\text{pure solute}) + G(\text{pure solvent})]ΔG(s) < 0: mixing is favourable
ΔG(s) > 0: mixing is unfavourable
ΔS(mixing) is always > 0, so unfavourable mixing comes from a large positive ΔH (loss of strong solute–solute or solvent–solvent interactions that are not replaced by solute–solvent interactions)
Intermolecular forces are the reason your rain jacket is waterproof (nonpolar polymer repels polar water) and why acetone removes nail polish (similar polarity, "like dissolves like"). The boiling-point trends you learn here explain why small alcohols evaporate quickly from your skin (low molecular weight, moderate hydrogen bonding) while larger oils do not. Pharmaceutical chemists tune solubility by adding or removing polar functional groups so that a drug dissolves in the bloodstream but also crosses nonpolar cell membranes.
Students often think "higher molecular weight always means higher boiling point." It does not. Polarisability and molecular shape matter too. CF₃CF₃ is heavier than CH₃CF₃ but boils lower because its symmetry kills the dipole and its fluorines reduce polarisability.
Students confuse intermolecular forces with intramolecular bonds. A C–O bond is intramolecular. A hydrogen bond between two alcohol molecules is intermolecular. Only IMFs break during boiling.
Students assume that if a molecule contains fluorine, it must form strong hydrogen bonds. Fluorine in C–F bonds is a poor hydrogen-bond acceptor because its lone pairs are tightly held. Only H–F (the bond, not C–F) participates meaningfully in hydrogen bonding.
Students forget that ΔS of mixing is always positive. Unfavourable mixing (ΔG > 0) comes from enthalpy, not entropy. If the mixed state has weaker intermolecular forces than the separated components, ΔH is large and positive, and that can outweigh the favourable TΔS term.
⚠️ Ranking boiling points of 3-4 molecules is a classic exam question. You will be given structures and asked to order them. Use the hierarchy: hydrogen bonding > dipole-dipole > London dispersion.
⚠️ Know the difference between alkyl, vinylic, and aryl halides cold. Exam questions often ask "classify this halide" as a first step before asking about reactivity.
⚠️ Be ready to explain why mixing is favourable or unfavourable using ΔG = ΔH – TΔS. Saying "like dissolves like" is not enough; they want the thermodynamic reasoning.
⚠️ Symmetry and polarisability are easy marks if you understand them, but easy to lose if you rely only on molecular weight.
True or false: London dispersion forces are only present in nonpolar molecules.
False. LDF are present in all molecules. They are simply the only IMF in nonpolar molecules.
Fill in the blank: A vinylic halide has its halogen bonded to an ____-hybridised carbon.
sp²
True or false: ΔS of mixing is always negative.
False. ΔS of mixing is always positive (mixing increases disorder).
Fill in the blank: Hydrogen bonding requires H bonded to ____, ____, or ____.
N, O, or F
True or false: A molecule with a higher molecular weight always has a higher boiling point than one with a lower molecular weight.
False. Symmetry, polarisability, and dipole moment can override molecular weight.
Q: Classify the following as alkyl, vinylic, or aryl halide: chlorobenzene; 2-bromopropane; vinyl chloride (CH₂=CHCl).
A: Chlorobenzene is an aryl halide (Cl on aromatic ring). 2-Bromopropane is an alkyl halide (Br on sp³ carbon). Vinyl chloride is a vinylic halide (Cl on sp² carbon of a C=C double bond).
Q: Rank the following in order of increasing boiling point: CH₃F, CH₃Cl, CH₃Br, CH₃I.
A: CH₃F < CH₃Cl < CH₃Br < CH₃I. As you go down the halogen group, molecular weight and polarisability increase, strengthening London dispersion forces. The dipole moment decreases going down the group, but LDF dominate here.
Q: Explain why ethanol (CH₃CH₂OH) is miscible with water but diethyl ether (CH₃CH₂OCH₂CH₃) is only slightly soluble.
A: Ethanol can both donate and accept hydrogen bonds with water (O–H group). Diethyl ether can accept hydrogen bonds (lone pairs on O) but cannot donate one (no O–H bond), so its interactions with water are weaker. It also has a larger nonpolar surface area.
Q: Two molecules have the same molecular formula but one is linear and the other is highly branched. Which has the higher boiling point, and why?
A: The linear isomer has the higher boiling point. Its extended shape provides more surface area for London dispersion interactions between neighbouring molecules. Branching reduces contact area and weakens LDF.
Q: Using ΔG = ΔH – TΔS, explain why hexane and water do not mix.
A: Water has strong hydrogen bonds (high solvent–solvent interactions). Hexane is nonpolar and cannot form hydrogen bonds with water. When mixed, the strong water–water interactions are lost without being replaced by comparably strong hexane–water interactions, so ΔH of mixing is large and positive. Even though ΔS of mixing is positive, the large positive ΔH makes ΔG positive, and mixing is unfavourable.
This material connects directly to Chapters 9 and 10 (substitution and elimination reactions), where the type of halide and the leaving-group ability of the halogen determine which reaction pathway dominates. The solubility principles here also reappear in biochemistry when you consider how drugs partition between aqueous (blood) and lipid (cell membrane) environments. The thermodynamic framework (ΔG = ΔH – TΔS) is the same one you will use in Chapter 6 for reaction equilibria.
Alkyl halide, haloalkane, vinylic halide, aryl halide, sp3 carbon, sp2 carbon, intermolecular forces, IMF, London dispersion forces, van der Waals forces, temporary dipoles, induced dipoles, polarisability, polarizability, dipole-dipole interaction, hydrogen bonding, H-bond, like dissolves like, miscible, solubility, Gibbs free energy, delta G, enthalpy of mixing, entropy of mixing, boiling point, molecular weight, surface area, branching, organic chemistry chapter 8, UW organic chemistry, Forrest Michael