Opening Image: Urea. The carbon, oxygen and nitrogen atoms are sp²-hybridized. The overlap of the adjacent and parallel p orbitals creates a π molecular orbital system. This has two consequences: it prevents rotation about the axes connecting the O–C–N nuclei, and it permits delocalization of the electron pair occupying a π orbital. Because rotation is prevented, the urea molecule is both planar and rigid.

The hydrogen bond is important to an understanding of molecular recognition. Hydrogen bonds provides no net free energy but are responsible for aligning atoms and holding them at precise distances and angles to each other in a noncovalent complex. Ionic interactions are also appealing because, as everyone knows, unlike charges attract each other. When a positive ion encounters a negative ion and a complex is formed, it is referred to as an ion pair. As we shall see, hydrogen bonding and ion pairing often occur together.

Urea dimer. Urea forms hydrogen-bonded dimers in the solid state and in organic solvents.

Urea is a classical example of hydrogen bonding to carbonyl groups. As mentioned above, urea is a planar molecule with sp²-hybridized carbon, oxygen, and nitrogen atoms. Thus, the urea molecule has six sites for hydrogen bonding: four donor N–H groups and two acceptor site, namely the lone pair electrons of the carbonyl oxygen atom. The oxygen sp² lone pair lobes are in the R₂C=O plane and form angles of about 120° with the C=O bond.

The two hydrogen bonds are linear (180°); furthermore, the N–H dipole points directly at the sp² lone pair orbitals of the oxygen atoms. In a given set of hydrogen bonds, the strongest hydrogen bonds are those closest to perfect geometry. And that is certainly the case for the two hydrogen bonds in the urea dimer.

Toggle spacefill.

Association Constants Are a Quantitative Measure of the Stability of Noncovalent Complexes

It's important to understand the stability of the urea dimer in the context of hydrogen bonding. Textbooks often list values for the enthalpy of different hydrogen bonds in the gas phase. However, in aqueous solutions the competition with water for hydrogen bonding sites weakens the free energy of hydrogen-bonded complexes significantly. This is not to say that hydrogen bonds are not important -- they are of paramount importance in molecular recognition. But this is something quite different from overall stability.

Let's consider the formation of urea dimers from a thermodynamic viewpoint. In water, the association constant is about 0.04 M^-1. In other words, the concentration of urea dimers is insignificant at concentrations below 10 M. However, the association constant in benzene is 40 M^-1, so urea exists only as the dimer in benzene. Why is the urea dimer a thousand times more stable in benzene than in water? The answer is that water forms hydrogen bonds to urea and at a concentration of 55 M, water molecules fill the urea hydrogen bonding sites most of the time.

Acetic acid dimer. Toggle spacefill.

Another example is acetic acid. The oxygen and carbon atoms in the –COOH group are sp² hybridized. The geometry for hydrogen bonding in the dimer is perfect and acetic acid forms a stable dimer in benzene. The dimer is so stable that it also exists in the vapor state. However, acetic acid dimers do not exist in aqueous solutions. In fact, the dissociation of acetic acid into carboxylate anions and H⁺ in dilute aqueous solutions is driven the strong hydrogen bonding of water to the dissociation products.

Guanidine Carries a Delocalized Positive Charge

Guanidinium ion. The guanidino group is found in many biomolecules, including the amino acid arginine. Toggle spacefilled model.

What a beautiful molecule -- it is planar, all C–N bonds are equal, and all bond angles are perfectly 120°. Based on these observations, describe the electronic structure of the guanidinium cation.

Show Answer

The carbon atom has only six electrons. It is sp² hybridized and is bonded to three nitrogen atoms, with bond angles of exactly 120°. The three nitrogen atoms are also sp² hybridized. The filled p orbitals can overlap with the empty p orbital of the carbon atom. Therefore, the positive charge is not confined to the carbon atom but is delocalized over the four sp² hybridized atoms. The ability to spread the positive charge over four atoms rather than concentrating it on just one atom accounts for the stability of the guanidinium ion, making it the "world's most stable carbocation."

Long Range Charge-charge Interactions Account for the Stability of Ion Pairs

Guanidine·acetate ion pair. This model is built from atomic coordinates for an arginine-glutamate ion pair found at the interface of a protein-protein complex. The positively charged guanidino group is found in the side chain of arginine, whereas the "acetate" anion is modeled after the side chain carboxyl of the glutamate residue in the protein. The interaction is entirely electostatic. Although there is no hydrogen bonding in this example, most ionic interactions involving functional groups found in biomolecules also involve hydrogen bonding. Toggle spacefilled model.

How strong is the electrostatic interaction? Answer.

Guanidinium Cations Can Also Serve as Proton Donors

Ideal guanidine·acetate ion pair. This model is based on energy minimization of the previous model. The strong hydrogen bonding in the ion pair accounts for the planar conformation of this ion pair. Toggle spacefilled model.

The association constant for the guanidinium acetate ion pair in aqueous solutions is much higher (about 150 times) than that for the urea dimer. The greater stability of the hydrogen-bonded complex is a consequence of additional long range coulombic charge-charge interactions.

Guanidinium cations are a prominent recognition motif for carboxylate and phosphate binding in nature. However, simple ion pairing based on guanidinium cations is not strong enough to achieve an efficient complexation in aqueous solvents. For example, the association constant for the guanidium acetate ion pair in water is 6 M^-1. Nature uses the less polar microenvironment of a protein to shield the ion pair from the solvent thereby increasing complex stability.

The Nucleic Acid Base Guanine Contains the Guanidine Center

Acetate ions bind strongly to guanine and do not bind to any other base. The association constant for carboxylate-guanine association is much higher (about 30 times) than that for guanine-cytosine base-pair formation. Carboxylate ions form two hydrogen bonds to the NH₂ and NH of guanine.

Guanine is the only base that is able to form these two hydrogen bonds with carboxylate ions; none of the other bases has two donor groups in the correct position. This explains the highly selective interaction of carboxylate ions with guanine.

Cooperative Interactions

Hydrogen bonds do not normally occur as isolated entities but form networks. Within these networks, hydrogen bond energies are not additive. Rather, the free energy is greater than the sum of the individual interactions. In other words, the system displays cooperativity. We'll return to this very important concept in later lectures. For now here's a glance of how nature uses networks of hydrogen bonding to achieve high specificity and affinity in molecular complexes and assemblies.

A charge-charge interaction between a phosphate ion and the guanidino group of an arginine side chain in a phosphate binding protein. The α-carbon atom of arginine is colored dark gray. Guanidium cations can also serve as proton donors. Toggle spacefill.

The 3D structure of proteins makes cooperative (synergistic) interactions possible. Looking at the same phosphate binding protein more closely, we see a hydrogen-bonded network of charge-charge interactions. Display the side chain of aspartate. Toggle hydrogen bonds.

The strong interaction between the side chain carboxyl and guanidino groups ensures that the guanidino group is precisely positioned to bind the phosphate ion.