Lab 2: Relative Basicity of Nitrogenous Bases
Acid-base reactions have been rationalized in many ways. Using Brønsted's simple definition, a good base is one that readily accepts a proton. Thus, a base stronger than ammonia would have an equilibrium constant greater than 1 for equation 1. Any theory of basicity can be quantified by measuring the equilibrium concentrations of the base and its conjugate acid or their relative stabilities. In this laboratory, we'll be comparing the gas-phase basicities of several nitrogenous bases relative to that of ammonia. Factors that stabilize the conjugate acid will make the base stronger but factors that stabilize the base itself will make it weaker. For example, OH- is a stronger base than NH3 for two reasons, because OH- is less stable than NH3 and because the conjugate acid, H2O, is more stable than the conjugate acid, NH4+. To put it another way, any factor that stabilizes the right side of equation 1 will strengthen the base relative to ammonia but factors that stabilize the left side will weaken it.
Looking at equation 1, if the base, B, is a nitrogenous base then the equation is considered an isodesmic reaction--one in which the number of each type of chemical bond is maintained on both sides of the chemical equation. The left side of the reaction has 4 N-H bonds and so does the right hand side. Isodesmic reactions are most accurate for computational chemistry because errors tend to cancel each other on both sides of the equation.
According to Lewis, a good base is one that can donate a pair of electrons well. Any resonance structure or hyperconjugative effect that decreases the availability of the lone pair will decrease its basicity. For this reason, phenolate (PhO-) is a weaker base than hydroxide (OH-) because a lone pair of electrons and the negative charge are delocalized into the benzene ring by resonance.
The most comprehensive definition of an acid/base interaction involves molecular orbitals. A good base according to the orbital definition, is a molecule that has a well-positioned and high-lying occupied orbital (usually the HOMO) to interact with an acid's low-lying unoccupied orbital (usually the LUMO). A Brønsted acid functions as an acid, in this model, because it has a low-lying unoccupied X-H antibonding orbital (e.g., the O-H antibonding orbitals in H3O+). The diagram below, shows the MO interaction of ammonia's HOMO with a Brønsted acid's empty X-H antibonding orbital. The closer in energy the two interacting orbitals are, the greater the stabilization of the occupied orbital. Thus, a base is generally stronger with a higher energy orbital and weaker with a lower energy orbital. Therefore, the MO explanation of why ammonia is a weaker base than hydroxide is that ammonia has a lower energy HOMO than hydroxide.
The enthalpy of the following gas phase reactions have been measured experimentally. This lab seeks to determine to what extent these values can be reproduced computationally and to use these computations to help explain the trends in basicity.
Experimental Relative Enthalpies for Equation 1 |
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The computational enthalpy is found using the following equations:
A minimum of three calculations will be necessary for each molecule (color coding in DH eq'n corresponds to calculation color coding below):
The AM1 level of theory is used for the geometry optimization and frequency calculations. Compare AM1 energies with classmates to ensure that the global minimum has been found then record the zero-point vibrational energy, ZPVE, and heat capacity, Cv, for each molecule. These numbers are found in the frequency log file: select
While the AM1 method is fast, it is not usually able to give relative energies that match experimental accuracy. Much slower methods can match experimental accuracy but take too long to be practical. A good compromise is to calculate the geometry at a lower level of theory and then use a higher level of theory to calculate the single-point energy at the lower-level geometry. Here, we will use the AM1-optimized geometry and calculate a "single point energy" using HF/6-31G(d). This is type of calculation is represented, in shorthand, as HF/6-31G(d)//AM1 and will provide us with a moderately reliable answer with a minimum demand on computational resources. When running the Hartree-Fock single point, make sure to save a checkpoint file in your folder so that you can visualize the orbitals later. To do this, select >calculate>Gaussian, change the checkpoint file under 
by clicking next to %chk= and selecting the
For each of the bases assigned to you, look for the molecular orbital in the base that is responsible for its basicity (i.e., the one most likely to interact with the transferring proton). In many cases this is the HOMO, but it is not necessarily so. Look for an orbital that is roughly where the new N-H bond is found in the conjugate acid. You need not visualize the orbitals of the conjugate acid. You can look at the orbitals and their energy levels in the MO window (shown below), by selecting
| Orbitals cannot be printed directly from the MO window. To print an orbital, select |
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Results Format:
Questions to think about for discussion section:
In your discussion section, make sure to consider all three methods of describing basicity (Brønsted, Lewis, MO).
Acetonitrile
Pyridine
Quinuclidine