Class Notes (1,100,000)
US (470,000)
UofL (900)
CHEM (50)
Lecture 3

CHEM 202 Lecture Notes - Lecture 3: Covalent Bond, Electronegativity, Vsepr Theory


Department
Chemistry
Course Code
CHEM 202
Professor
Neal Stolowich
Lecture
3

This preview shows page 1. to view the full 5 pages of the document.
Lecture 3
Chap 37 – Intermolecular Forces
Intermolecular Forces– An introduction
We’ll start off by introducing all the intermolecular forces here. Recall intermolecular forces involve
the interaction between molecules, while intramolecular forces are the bonding forces within a
molecule (and are typically orders of magnitude stronger: hundreds to thousands kJ/mol)
ion – dipole: strongest- involves full charge of an ion attracting dipole charge, therefore is also the
strongest intermolecular force (chap 39)
hydrogen bond: special dipole. Involves H and O, N, and F bonds usually
dipole – dipole: not quite as strong as ion-dipole since on partial charges are present.
Ion/Dipole Induced dipoles: certain functional groups can polarize neighboring molecules, thus
inducing a dipole. polarizability increases with radii (electrons farther away from nucleus
is easier to polarize!) decreases across period, increases down group (Chap 39)
Dispersive forces – weakest of the intermolecular forces, involves London forces associated with
non-polar interactions.
These are summarized in the Silberberg figure (IFcomparo)
Recall/review from Chem 201 what dipoles and polar compounds are.
Consider water: the O – H bond is a polar or polarized bond due to the electronegativity
difference between O and H. The electrons in the covalent bond are pulled towards the oxygen
resulting in a partial positive charge on oxygen, and a partial positive charge on hydrogen. Due to the
overall geometry of the molecule (AX2E2; tetrahedral class) the water molecule is polar overall as well.
Carbon dioxide on the other hand – while containing polar C=O bonds, due to symmetry (AX2; linear)
the dipoles cancel out and the molecule is consider non-polar
Chapter 37 will first consider three primary types for single, pure substances (the other two will be
considered for mixtures, chap 39) - dipole-dipole, H bonding, and dispersion (37.1)
A. Dipole-dipole (37.2) : Review what is a dipole (chap 29). Many molecules contain dipoles, and
dipole-dipole IF is available to all polar molecules. These will be all examples of permanent
dipoles.
Recall VSEPR geometry dictates whether a molecule containing dipoles is a polar molecule or not,
and this is important for considering whether dipole-dipole IFs occur.
H2O, CH3F, ClF3, ClF5 and SF4 all have dipoles, are polar molecules, and experience dipole-dipole
attraction in their solid and liquid phases.
find more resources at oneclass.com
find more resources at oneclass.com
You're Reading a Preview

Unlock to view full version

Only page 1 are available for preview. Some parts have been intentionally blurred.

BeH2, BH3, SiCl4, XeF2, and AsCl5 may contain dipolar bonds, however are non-polar molecules
and lack dipole-dipole attraction in any phase.
Dipole-dipole interactions are generally weak (5 – 25 kJ/mol) as they involve electrostatic attraction
between the opposite partial charges within the dipole. The more polar the molecule, the stronger
the dipole-dipole IFs will be.
B. Hydrogen bonding (37.3): Hydrogen bonding is an attraction between a hydrogen atom of a
very polar bond and a nearby orbital, typically containing lone pair electrons available for
interaction. O – H, N – H, and F – H are common very polar bonds containing H, and also have
the lone pairs available to complete the hydrogen bond. (see figures, p 380) The more polar the
bond, the stronger the H-bond will be. Hydrogen bonding is a very important interaction. If
present in small molecules, it will be the strongest IF. Even in large molecules, extensive use of H-
bonding will dominant other IFs. Such can be the case in RNA, DNA and proteins.
Water is the ultimate hydrogen bonding substance. It has two O – H bonds and two lone pair
orbitals forming the ultimate H-bonding network (p381). This is what gives water all of its unique
properties, such as high (relatively) melting and boiling points, as well as the increase in density
upon melting and the phase diagram described earlier.
Notes to keep in mind: while we call it H-bonding – which infers a bond – no full chemical bond is
formed. Secondly, while H-bonding may look like a dipole-dipole attraction, it is not – as seen
above, dipole-dipole attraction depends on molecular polarity, H-bonding does not! (ie, boric acid
is non-polar in the solid phase, but has strong H-bonds to its neighbors)
C. Dispersion (37.4) – dispersive IFs are generally the weakest, however can “accumulate” within
a molecule and therefore become quite important. Dispersive IF’s typically occur in non-polar
molecules, such as hydrocarbons. Dispersion is not straightforward and can involve two types of
temporary dipoles: an induced dipole and an instantaneous dipole.
An induced dipole only exists when a neighbor causes it. Typically, an induced dipole is a result of
distortion of the electron orbitals. Bringing a dipole near a non-polar molecule or atom can distort
the electrons, inducing a temporary dipole in the target. And the more polarizable a molecule or
atom is, the greater the induced dipole, and stronger the dispersion can be.
find more resources at oneclass.com
find more resources at oneclass.com
You're Reading a Preview

Unlock to view full version