Using Molecular Structure to Predict Equilibrium

Electronegativity is the tendency of an atom to attract electrons toward itself when forming bonds. Atoms with increased electronegativity are more stable after the addition of a negative charge. One example of the effect of electronegativity on acidity is demonstrated by the comparison of a hydrogen connected to an oxygen atom versus a hydrogen connected to a carbon atom. Oxygen is the more electronegative atom, and so it is more stable with a negative charge than a carbon with a negative charge. This stability results in a lower pK_a value for oxygen than carbon, which indicates that oxygen is a stronger acid. In general, molecules with a hydrogen bonded to oxygen, such as methanol, are more acidic than molecules with a hydrogen bonded to carbon, such as ethane. Electronegativity increases from left to right across a period (row) on the periodic table and from bottom to top within a group (column) on the periodic table. Because electronegativity is a periodic trend, the atomic effect of electronegativity can be applied using a periodic trend. Moving from left to right across the periodic table within a period increases the acidity. Likewise, moving from bottom to top within a group increases the acidity.
The more stable and weaker a conjugate base, the stronger the acid that generated that conjugate base. Fluorine is the most electronegative atom. When HF dissociates, it forms fluoride, F^–, which is very stable because fluoride is more electronegative than oxygen, nitrogen, and carbon. The size of anions affects their stability in a similar fashion to resonance delocalization. Anions that are larger have more room to spread out a negative charge and are therefore more stable. This is what makes hydrogen iodide $({\rm {p}}K_{\rm {a}}=-10)$ a stronger acid than hydrogen fluoride $({\rm {p}}K_{\rm {a}}=3.2)$.
Size increases down a group (column) in the periodic table; therefore, the strength of an acid increases down a column in the periodic table. In group 17, the halogens, the size increases down the chart; so iodine is the largest atom and, thus, HI is the strongest HX acid. Fluorine is the smallest atom, and HF is the weakest HX acid. Larger atoms are more polarizable and have a much larger volume. Thus, larger atoms are able to disperse the negative charge that results from the dissociation of HX better than smaller atoms. The strongest acids have the most stable conjugate base. HI is a very strong acid. When HI dissociates, iodide is a very stable and weak conjugate base. Because it is so large, it is able to spread the negative charge over a very large volume and form a stable, weak conjugate base.

It may seem counterintuitive that HI is a stronger acid than HF, because fluorine is more electronegative than iodine. But iodine is larger than fluorine. Therefore, there are two competing trends with the halogens—size and electronegativity. Size is a more important trend than electronegativity for the halogens, and HI is the stronger acid.

Delocalization of a negative charge is possible on molecules that have resonance structures. A resonance structure is one of two or more Lewis structures with multiple equivalent representations. This delocalization of the negative charge is a stabilizing factor, and the more atoms the charge is spread over, the more stable the anion. Anion stability equates to base stability, which means a stable base is a weak base. Therefore, if a base has resonance structures, it will be a stable, weak base and is generated from a strong acid. For example, the carboxylate anion is stabilized by resonance, while the ethoxide anion is not. Resonance is the method that organic chemists use to deal with organic molecules that have two or more Lewis structures with multiple equivalent representations. When a conjugate base is more stable, it is weaker. These conjugate bases come from acids that dissociate,and the equilibrium shifts toward the products (conjugate base). Resonance structures in a conjugate base make the molecule a weaker and more stable conjugate base, which means it was formed from a stronger acid. Resonance will spread out or delocalize the negative charge over multiple atoms. This is similar to how a large iodide spreads out the negative charge over its larger volume. Without resonance, a charge is "locked in"—localized on one atom—which makes that structure a less stable and stronger conjugate base. Phenol has resonance structures, but cyclohexanol has none. Therefore, phenol is a stronger acid because its conjugate base is stabilized by the resonance structures it forms. Resonance structures help stabilize a conjugate base and, thus, decrease the basicity of the conjugate base, which means that the conjugate base was generated from a stronger acid.

An electronegativity difference between atoms creates a bond polarity, which leads to the inductive effect. The inductive effect relies on the intrinsic tendency of the molecule to withdraw or release electrons because of its electronegativity. As the distance between the electronegative atom and the acidic hydrogen atom increases, the inductive effect weakens. Electronegativity increases from left to right across a row (period) on the periodic table and from bottom to top within a group (column) on the periodic table. Fluorine, chlorine, nitrogen, and oxygen are some atoms that often create an inductive effect within a molecule.

The inductive effect weakens conjugate bases. Unlike the resonance effect, double bonds are not necessary, but the polarity of the bond is essential. The inductive effect is a permanent displacement of a shared electron pair of a covalent bond towards the more electronegative atoms or groups. As the inductive effect withdraws the negative charge from an atom, the charge is split over a larger volume and stabilizes. The inductive effect is similar to the polarizability of the electron cloud of larger atoms.

The inductive effect also stabilizes the anion, but to a lesser extent than resonance delocalization. The difference in electronegativities between two atoms causes the flow of electrons from the less electronegative atom toward the more electronegative atom. In a simple molecule such as water (H₂O), this effect yields a bond dipole with a slightly positive end and a slightly negative end. In larger, more complex molecules, there might be many such bond dipoles. For example, there is a large difference in acidity between acetic acid and trichloroacetic acid. This difference is because carbon is more electronegative than hydrogen. In a carbon-hydrogen bond, the more electronegative carbon pulls the electron pair towards it and away from the hydrogen. Therefore, the carbon is on the negative end of carbon-hydrogen dipole with the hydrogen being on the positive end of the dipole. When the hydrogens are replaced with chlorines, the bond dipoles are reversed. This now puts carbon on the positive end of the carbon-hydrogen dipole, and the inductive effect works beneficially to stabilize the negative charge on the carboxylate anion. Acetic acid (CH₃COOH) is acidic because of both the atom and resonance effect. The acidic proton of acetic acid is connected to an electronegative oxygen atom, and the oxygen atom attached to the acidic proton is adjacent to a carbon-oxygen double bond. Because of this, acetic acid has a pK_a of 4.76. Both acetic acid and trifluoroacetic acid (CF₃COOH) are acidic because of both the atom and the resonance effect. However, trifluoroacetic acid also has a very strong inductive effect because of the presence of the three fluorine atoms and is, therefore, more acidic than acetic acid. Additionally, acetic acid contains a methyl group, which is an electron donating group (as opposed to halogens which are electron withdrawing groups). Electron donating groups create a reverse inductive effect and make an acid a weaker acid. Trifluoroacetic acid has a pK_a of 0.23.