Hydrogen bonds are a favorite of the MCAT! So understand them well, and you will be able to answer every question pertaining to their nature, behavior, and effects. A hydrogen bond is actually only a specific, unusually strong form of dipole–dipole interaction, which may be intra- or intermolecular. Please understand that hydrogen bonds are not actually bonds—there is no sharing or transfer of electrons between two atoms. When hydrogen is bound to a highly electronegative atom, such as fluorine, oxygen, or nitrogen, the hydrogen atom carries little of the electron density of the covalent bond. The hydrogen atom acts essentially as a naked proton (it is a little scandalous!). The positively charged hydrogen atom interacts with the partial negative charge of oxygen, nitrogen, or fluorine on nearby molecules. Substances that display hydrogen bonding tend to have unusually high boiling points compared with compounds of similar molecular formula that do not hydrogen bond. The difference derives from the energy required to break the hydrogen bonds. A very good example of this difference, which you absolutely need to understand for the MCAT, is the difference between a carboxylic acid, which can form hydrogen bonds between molecules of itself, and the acyl halide, a derivative of the carboxylic acid, which has substituted a halogen for the hydroxyl group and which cannot form hydrogen bonds with molecules of itself. The boiling point of propanoic acid (MW = 74 g/mol) is 141°C, but the boiling point of propanoyl chloride (MW = 92.5 g/mol) is 80°C. Hydrogen bonding is particularly important in the behavior of water, alcohols, amines, and carboxylic acids. It is no overstatement to say that if it weren’t for water’s ability to form hydrogen bonds and exist in the liquid state at room temperature, we would not exist (at least not in the form we recognize as “human”).
Conclusion
This chapter built upon our knowledge of the atom and the trends demonstrated by the elements in the periodic table to understand the different ways by which atoms partner together to form compounds, either (a) exchanging electrons to form ions, which are then held together by the electrostatic force of attraction that exists between opposite charges, or (b) sharing electrons to form covalent bonds. We discussed the nature and characteristics of covalent bonds, noting their relative lengths and energies as well as polarities. The review of Lewis structures and VSEPR theory will prepare you for predicting likely bond arrangements, resonance structures, and molecular geometry. Finally, we compared the relative strengths of the most important intermolecular electrostatic interactions, noting that even the strongest of these, the hydrogen bond, is still much weaker than an actual covalent bond. The next time you’re searing meat in a skillet or browning toast in a toaster, take a moment to consider what’s actually going on at the atomic and molecular level. It’s not just cooking; it’s science!
CONCEPTS TO REMEMBER
Atoms come together to form compounds through bonds. Many atoms interact with other atoms in order to achieve the octet electron configuration in their valence shell. There are two types of bonds: ionic and covalent.
Ionic bonds form when a very electronegative atom gains one or more electrons from a much less electronegative atom; the strong electrostatic force of attraction between opposite charges holds the resulting anion and cation together. Covalent bonds form when two atoms share one or more pairs of electrons. The force of the bond is due to the electrostatic attraction between the electrons in the bond and each of the positively charged nuclei of the bonding atoms.
Covalent bonds are characterized by their length, energy, and polarity. A single covalent bond exists when two atoms share one pair of electrons, a double bond exists when two atoms share two pairs of electrons, and a triple bond exists when two atoms share three pairs of electrons. Triple bonds are the shortest and have the highest bond energy. Single bonds are the longest and have the lowest bond energy.
Although all covalent bonds involve a sharing of one or more electron pairs, the electron density of the bonding electrons may not be distributed equally between the bonding atoms. This gives rise to a polar bond, with the more electronegative atom receiving the larger share of the bonding electron density.
Lewis structures are bookkeeping devices used to keep track of the valence electrons of atoms and their various possible arrangements in a compound. Evaluation of formal charge on each atom in the arrangement can help determine which arrangement is most likely to be representative of the actual bond connectivity. Lewis structures that demonstrate the same atomic arrangement but differ in the distribution of electron pairs are called resonance structures. The actual molecular positioning of electrons is a “weighted” hybrid of all the possible resonance structures called a resonance hybrid.
VSEPR theory helps us predict the actual three-dimensional arrangement of bonded and nonbonded electron pairs around the central atom in a compound. The theory states that nonbonded electrons (lone pairs) exert strong electrostatic repulsive forces against the bonded pairs of electrons and, as a result, the electron pairs arrange themselves as far apart as possible in order to minimize the repulsive forces.
A molecular dipole is the resultant of all the bond dipole vectors in the molecule. A compound consisting of only nonpolar bonds will by definition be nonpolar. A compound consisting of one polar bond will by definition be polar (although perhaps insignificantly). A compound consisting of two or more polar bonds may be polar or nonpolar, depending on the molecular geometry and the vector addition of the bond dipoles.
Atomic orbitals overlap their + and - lobes to result in bonding or antibonding molecular orbitals. When the signs of the overlapping atomic orbitals are the same, the result is a bonding molecular orbital. When the signs are opposite, the result is an antibonding orbital.
Sigma bonds involve head-to-head (or end-to-end) overlap of electron cloud density and thus allow for relatively low-energy rotation. Pi bonds involve parallel overlap of electron cloud density and thus do not allow for low-energy rotation.
The intermolecular interactions are electrostatic interactions that weakly hold molecules together. They generally are significant only over short distances and in the solid or liquid state. The weakest intermolecular interaction is the dispersion force (London force), which results from the moment-by-moment changing unequal distribution of the electron cloud density across molecules. A more moderate-strength intermolecular interaction is the dipole–dipole interaction, which exists between the opposite ends of permanent molecular dipoles. The strongest of these intermolecular interactions is the hydrogen bond, which is a special case of the dipole–dipole interaction between the hydrogen atom attached to either oxygen, nitrogen, or fluorine and another oxygen, nitrogen, or fluorine on another molecule—or sometimes, if the molecule is large enough, even the same molecule. One of the most important occurrences of hydrogen bonding is that between water molecules.
EQUATIONS TO REMEMBER
µ = qr
Formal charge = V - Nnonbonding -½Nbonding
Practice Questions
1. What is the character of the bond in carbon monoxide?