Chemistry · Ch 10 — States of Matter

Types of Intermolecular Forces

10.2.1

Types of Intermolecular Forces

Polar molecules -- those in which bonded atoms differ enough in electronegativity that the shared electrons sit closer to one atom than the other -- carry a permanent separation of positive and negative charge called a dipole. This is measured by the dipole moment, $\mu = Q \times r$ , the product of the magnitude of charge $Q$ and the distance $r$ between the centres of positive and negative charge, expressed in debye (D); it is a vector, drawn as an arrow with its tail at the positive centre and its head at the negative centre (e.g. in HCl or HF). Four kinds of weak intermolecular attraction act between such molecules. (i) Dipole-dipole interactions arise when the positive end of one polar molecule is attracted to the negative end of a neighbouring polar molecule; these are weak (about $3$ -- $4$ kJ mol $^{-1}$ ) and matter mainly in the solid and liquid states, and stronger dipole moments generally raise the boiling point (e.g. CH3CN, with the largest dipole moment among molecules of similar molar mass, has the highest boiling point). (ii) Ion-dipole interactions act between a full ion (cation or anion) and the partial charges of a nearby polar molecule, e.g. between Na $^+$ and the negative end of H2O; the smaller and more highly charged the ion, the stronger this attraction, so the order is Na $^+$ < Mg $^{2+}$ < Al $^{3+}$ . (iii) Dipole-induced dipole interaction occurs when a polar molecule temporarily distorts the electron cloud of a nearby non-polar molecule, inducing a short-lived dipole in it. (iv) London dispersion forces -- also called van der Waals forces -- are the weakest of all and are the only attraction possible between non-polar molecules or noble-gas atoms; at any instant the electron cloud of an atom can be momentarily lopsided, creating a fleeting dipole that induces a matching fleeting dipole in a neighbour. Their strength grows with molecular size, mass, electron count and polarizability; a longer, less-compact chain (n-pentane, b.p. 309.4 K) has stronger dispersion forces than a compact spherical isomer of the same formula (neo-pentane, b.p. 282.7 K). (v) Hydrogen bonding is a special, comparatively strong dipole-dipole attraction (up to about $40$ kJ mol $^{-1}$ ) that forms when a hydrogen atom bonded to nitrogen, oxygen or fluorine is attracted to a lone pair on a nearby electronegative atom; …

Figure Fig 10.2(a)Fig. 10.2(a): A polar molecule (HCl)

What this figure shows. Shows the covalent HCl molecule with unequal sharing of the bonding electron pair between hydrogen and chlorine: a Greek delta (small-delta) with a plus sign marks a partial positive charge on the hydrogen atom, and a delta with a minus sign marks a partial negative charge on the chlorine atom (which has the greater electron density/electronegativity), with a small arrow (crossed at the tail) drawn from the positive to the negative centre to represent the dipole moment vecto …

Figure Fig 10.2(b)Fig. 10.2(b): Dipole-dipole interaction in liquid ICl

What this figure shows. Depicts several ICl molecules, each individually polar with a partial positive charge on iodine and partial negative charge on chlorine, oriented in the liquid so that the positive end of one ICl molecule lines up close to and is attracted toward the negative end of a neighbouring ICl molecule -- illustrating how many individual dipoles line up head-to-tail in a real dipolar liquid, the physical origin of dipole-dipole intermolecular …

Figure Fig 10.2(c)Fig. 10.2(c): Na+ ion - H2O interaction (hydrated Na+ ion)

What this figure shows. Shows a single, small, positively charged sodium ion surrounded by several water molecules, each oriented so that its partial-negative oxygen end points toward the central sodium cation, since opposite charges attract; this arrangement -- several polar solvent molecules clustering around and orienting toward a central ion -- is what is meant by 'hydration' of the ion, and is the visual example given for an ion-dipole interac …

Figure Fig 10.2(d)Fig. 10.2(d): Dipole-induced dipole interaction

What this figure shows. Shows, side by side, a polar molecule (with a permanent dipole, drawn with partial-charge labels) approaching a non-polar molecule; a further panel shows the same non-polar molecule now with an induced dipole -- its electron cloud has been momentarily pulled/distorted toward the polar molecule's positive end, generating a temporary partial charge separation in the previously non-polar molecule, which is then weakly attracted to the …

Figure Fig 10.3Fig. 10.3(a)/(b): n-pentane vs neo-pentane -- shape and boiling point

What this figure shows. Two structural sketches of the same molecular formula C5H12: panel (a) shows n-pentane as a longer, less compact, zig-zag chain of five carbon atoms (boiling point 309.4 K); panel (b) shows neo-pentane as a more compact, roughly spherical branched structure with a central carbon bonded to four methyl groups (boiling point 282.7 K, noticeably lower). The figure visually contrasts how a longer, more spread-out molecular shape allows greater surface contact between neighbouring molecules, and therefore stronger London dispersion forces and a higher boiling point, than a compact spherical isomer of identica …

Table Table 10.2Table 10.2: Effect of dipole moment on boiling point

Substance | Molar Mass (amu) | Dipole Moment (D) | Boiling point (K)

CH3-CH2-CH3 (propane) | 44.10 | 0.1 | 231

CH3-O-CH3 (dimethyl ether) | 46.07 | 1.3 | 248

CH3-Cl (chloromethane) | 50.49 | 1.9 | 249

CH3-CN (acetonitrile) | 41.05 | 3.9 | 355 …

Figure Fig 10.4Fig. 10.4: Hydrogen bonding in H2O and NH3

What this figure shows. Two small diagrams, one above the other. The upper one shows several H2O molecules linked together, with a solid line used for each molecule's own internal O-H covalent bonds and a dotted line drawn between the hydrogen of one water molecule and the oxygen (lone pair) of a neighbouring water molecule, representing the intermolecular hydrogen bond. The lower diagram shows the same pattern for NH3 molecules, with dotted lines connecting an N-H hydrogen on one ammonia molecule to the lone pair on the …

Figure Fig 10.5Fig. 10.5: Three-dimensional hydrogen-bond network in liquid water

What this figure shows. Shows many water molecules linked to each other in a sprawling three-dimensional lattice-like web via dotted hydrogen-bond lines, illustrating that because each H2O molecule has two hydrogen atoms available to donate and two oxygen lone pairs available to accept, every water molecule can simultaneously hydrogen-bond to up to four neighbours, building an extensive interconnected network rather than simple isolated pairs. …

Table Table 10.3Table 10.3: Comparison of intermolecular forces

Force | Strength | Characteristics

Ion-dipole | Moderate (10-50 kJ/mol) | Occurs between ions and polar solvents

Dipole-dipole | Weak (3-4 kJ/mol) | Occurs between polar molecules

London dispersion | Weak (1-10 kJ/mol) | Occurs between all molecules; strength depends on size and polarizability

Hydrogen bond | Moderate (10-40 kJ/mol) | Occurs between molecules with O-H, N-H, and H-F bonds …

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