Organic Chemistry I | University of Minnesota Twin Cities Source: Lecture notes (handwritten)
Tags: polarity, electronegativity, intermolecular forces, IMF, dipole, bond polarity, organic chemistry
Difficulty: Introductory
Prerequisites: General chemistry foundations, Lewis structures, VSEPR theory and molecular geometry.
This topic sits at the bridge between general chemistry and organic chemistry. You need to understand how electrons are shared (or not shared) between atoms before you can predict how molecules behave around each other. Bond polarity and intermolecular forces determine physical properties like boiling point, solubility, and viscosity, which come up constantly in organic chemistry. If you are comfortable with Lewis structures and VSEPR shapes, you are ready for this material.
Electronegativity differences (ΔEN) between bonded atoms tell you whether a bond is nonpolar covalent, polar covalent, or ionic. Whether the whole molecule is polar depends on both the bond polarities and the molecular shape. Intermolecular forces (ion-ion, dipole-dipole, dispersion) arise from these charge interactions and govern a molecule's physical properties.
Electronegativity (EN)
The tendency of an atom to attract shared electrons toward itself in a covalent bond. Think of it as how greedy an atom is for electrons.
Electronegativity difference (ΔEN)
The numerical difference in electronegativity between two bonded atoms, used to classify the bond type. In simple terms, ΔEN is a quick way to estimate whether a bond is nonpolar, polar, or ionic.
Nonpolar covalent bond
A bond where ΔEN < 0.4, meaning electrons are shared roughly equally between atoms. Think of it as a perfectly fair tug-of-war.
Polar covalent bond
A bond where ΔEN is between 0.4 and 1.9. The electrons are shared unevenly, creating partial positive (δ+) and partial negative (δ−) charges. In simple terms, one atom pulls harder, but not hard enough to steal the electrons entirely.
Ionic bond
A bond where ΔEN ≥ 1.9. The electronegativity difference is so large that one atom essentially takes the electron(s), producing full positive and negative ions.
Dipole
A separation of partial positive and partial negative charge within a bond or molecule. Think of it as one end being slightly positive and the other slightly negative, like a tiny bar magnet.
Molecular polarity
Whether the molecule as a whole has a net dipole moment. Depends on both bond polarities and molecular geometry.
Intermolecular forces (IMF)
The forces of attraction (or repulsion) between molecules, as opposed to the bonds within a molecule. In simple terms, IMFs are what hold separate molecules near each other in a liquid or solid.
Ion-ion interactions
The strongest type of intermolecular force, occurring between species with full positive and full negative charges.
Dipole-dipole interactions
Attractive forces between the partial charges of polar molecules. Weaker than ion-ion, but significant in polar substances.
London dispersion forces
Weak, temporary attractions that arise from momentary fluctuations in electron density. Present in all molecules (aka van der Waals forces, temporary dipoles, induced dipoles).
The electronegativity difference between two bonded atoms determines what type of bond forms.
ΔEN < 0.4 — Nonpolar covalent bond. Electrons are shared equally or nearly equally.
Example: H–H (ΔEN = 0). The electron cloud is symmetrically distributed around both hydrogen atoms.
ΔEN between 0.4 and 1.9 — Polar covalent bond. Electrons are shared unequally, creating partial charges (δ+ and δ−).
Example: H–Cl. Chlorine is more electronegative, so it pulls electron density toward itself. The H end is δ+ and the Cl end is δ−.
ΔEN ≥ 1.9 — Ionic bond. The difference is large enough that one atom effectively takes the electron, creating full ions.
Example: Li⁺ Cl⁻. Lithium loses its electron entirely to chlorine.
These are guidelines, not rigid cutoffs. Some textbooks use slightly different thresholds, but the principle is the same: larger ΔEN means more unequal sharing.
A molecule can contain polar bonds yet still be nonpolar overall. The key is molecular geometry.
Polar bonds in a molecule can make the molecule overall polar, unless the bond dipoles cancel each other out.
Cancellation happens when:
Two identical polar bonds point in exactly opposite directions (e.g. CO₂, which is linear, so the two C=O dipoles cancel).
Three or more identical polar bonds are arranged in a symmetric shape, such as trigonal planar (e.g. BF₃) or tetrahedral with all identical substituents (e.g. CF₄).
The molecule is polar when:
The geometry is asymmetric, so dipoles do not cancel.
Example: SO₂ is bent, so the two S=O dipoles do not cancel, and the molecule is polar.
Example: NH₃ is trigonal pyramidal (not trigonal planar), so the three N–H dipoles do not cancel.
Example: CHF₃ is tetrahedral but with mixed substituents (one H, three F), so the dipoles do not cancel perfectly.
The takeaway: you must analyse the shape of the molecule, not just its bonds, to determine overall polarity.
Worked examples from the lecture notes:
O=C=O (CO₂) — linear, symmetric — nonpolar
O=S with lone pairs (SO₂) — bent — polar
H₂C=CH₂ (ethylene) — symmetric planar — nonpolar
NH₃ — trigonal pyramidal — polar
F₂C=CF₂ — symmetric planar — nonpolar
H₂C=CF₂ — asymmetric substituents — polar
CF₄ — symmetric tetrahedral — nonpolar
CHF₃ — tetrahedral, mixed substituents — polar
Charge interactions between molecules create intermolecular forces. These are listed from strongest to weakest.
Ion-ion interactions — The strongest IMF. Occurs between species carrying full charges (e.g. Na⁺ and Cl⁻). These are the forces holding ionic compounds together in a crystal lattice.
Dipole-dipole interactions — Strong, but weaker than ion-ion. Occurs between the partial charges (δ+ and δ−) of polar molecules. The positive end of one molecule attracts the negative end of another.
London dispersion forces — The weakest individually, but omnipresent. Arises from temporary, instantaneous dipoles caused by random fluctuations in electron density. Every molecule experiences dispersion forces, whether polar or not. Larger molecules with more electrons have stronger dispersion forces.
In practice, most organic molecules experience a combination of these forces. A polar organic molecule will have both dipole-dipole and dispersion forces acting on it.
Intermolecular forces explain why water (polar, strong hydrogen bonding) has a surprisingly high boiling point for such a small molecule, while methane (nonpolar, dispersion only) is a gas at room temperature. This same principle is why oil and water do not mix: nonpolar oil molecules interact mainly through dispersion forces and have no affinity for the strong dipole interactions of water. In drug design, understanding molecular polarity helps predict whether a compound will dissolve in blood (polar) or pass through cell membranes (nonpolar interior).
Students often think that any molecule with polar bonds must be polar. It does not follow. If the geometry is symmetric (e.g. CO₂, CF₄), the bond dipoles cancel and the molecule is nonpolar.
Students sometimes confuse intermolecular forces with intramolecular bonds. IMFs act between separate molecules; covalent and ionic bonds act within a molecule or compound. Breaking IMFs is what happens when you boil a liquid, not breaking covalent bonds.
Dispersion forces are sometimes dismissed as negligible. They are individually weak, but they are the only force acting on nonpolar molecules and they increase substantially with molecular size. In large nonpolar molecules (like long-chain hydrocarbons), dispersion forces are strong enough to make the substance a solid at room temperature.
Students often assume ionic bonds and ion-ion intermolecular forces are the same thing. Ionic bonds hold the compound together internally; ion-ion IMFs describe the interactions between separate ionic species, for instance in solution.
⚠️ You will almost certainly be asked to determine whether a molecule is polar or nonpolar given its structure. The method: draw the geometry, identify the polar bonds, then check whether the dipoles cancel.
⚠️ Ranking intermolecular forces by strength (ion-ion > dipole-dipole > dispersion) is a standard exam question. Be ready to explain why, not just recite the order.
⚠️ Expect questions that give you two molecules and ask which has the higher boiling point. The one with stronger IMFs wins. If both are nonpolar, the larger molecule (more electrons, stronger dispersion) has the higher boiling point.
⚠️ ΔEN threshold values (0.4 and 1.9) are frequently tested. Know which range corresponds to which bond type.
True or false: CO₂ is a nonpolar molecule even though it contains polar C=O bonds.
True. The linear geometry means the two dipoles cancel.
True or false: Dispersion forces only occur in nonpolar molecules.
False. Dispersion forces are present in all molecules, polar or nonpolar.
Fill in the blank: A bond with ΔEN of 1.2 is classified as a ______ bond.
Polar covalent (ΔEN is between 0.4 and 1.9).
True or false: Ion-ion forces are weaker than dipole-dipole forces.
False. Ion-ion forces are the strongest type of intermolecular force.
Fill in the blank: To determine whether a molecule is polar, you need to know both the bond polarities and the molecular ______.
Geometry (or shape).
Q: Classify the bond in HF as nonpolar covalent, polar covalent, or ionic. Explain your reasoning.
A: Polar covalent. The ΔEN between H (2.2) and F (4.0) is 1.8, which falls between 0.4 and 1.9.
Q: Is CCl₄ a polar or nonpolar molecule? Why?
A: Nonpolar. Although each C–Cl bond is polar, CCl₄ has a tetrahedral geometry with four identical substituents, so the bond dipoles cancel symmetrically.
Q: Rank the following intermolecular forces from weakest to strongest: dipole-dipole, ion-ion, London dispersion.
A: London dispersion < dipole-dipole < ion-ion.
Q: Explain why NH₃ is polar but BF₃ is not, even though both contain polar bonds.
A: NH₃ is trigonal pyramidal (due to a lone pair on nitrogen), so the three N–H bond dipoles do not cancel. BF₃ is trigonal planar with 120° between each B–F bond, so the three dipoles cancel symmetrically.
Q: Two molecules have similar molecular weights. Molecule A is polar and molecule B is nonpolar. Which is likely to have the higher boiling point, and why?
A: Molecule A. It experiences dipole-dipole interactions in addition to dispersion forces, while molecule B relies on dispersion forces alone. Stronger total IMFs mean more energy is needed to boil the substance.
This connects directly to VSEPR theory and molecular geometry, because you need the 3D shape to determine whether dipoles cancel. If molecular geometry is fuzzy, revisit that material first.
It also connects to physical properties of organic compounds later in the course. Boiling points, melting points, and solubility are all governed by IMFs, and exam questions on those topics assume you can identify the dominant intermolecular forces.
In later chapters on reactions and mechanisms, understanding polarity helps predict where nucleophiles and electrophiles will attack, since reactive sites are often the most electron-rich or electron-poor parts of a molecule.
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