Section 3: Structural Considerations in Acid-Base Chemistry

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I am sure at this point, one is questioning, what are the properties of a strong acid? Structurally how can one differentiate between a strong acid and a weak acid? To answer these questions, we need to draw upon concepts discussed previously: resonance, polarizability, thermodynamics, electronegativity, and inductive effects. A thorough understanding of bonding is essential to explain specific trends.

Trend 1: Binary Acids

A binary acid has the form H – A where A is a nonmetal atom. Let’s consider the pK_a values across the second row of the periodic table.

Acid CH₄ NH₃ H₂O HF

pK_a 51 35 14 4

Conjugate Base CH₃^- NH₂^- OH^- F^-

Recall as pK_a decreases, acidity increases. Therefore, CH₄ is the weakest acid and HF is the strongest acid in this trend. To explain this trend, we need to consider the structure of the conjugate bases. Recall equilibrium and thermodynamics always favors the production of a weaker base. If we compare the structure of methide (CH₃^-) and fluoride (F^-) we observe a lone pair on carbon vs. a lone pair on fluorine. Using electronegativity, we can understand that the fluoride (electronegativity value = 4.0 units) will attract (hold onto) and stabilize the lone pair more readily than carbon (electronegativity value = 2.5 units). Because the fluoride attracts the lone pair more readily, it is less available to react or accept a proton than the lone pair on carbon. We can use this rationale as explain why water is a stronger acid than ammonia.

While electronegativity can be useful in rationalizing binary acid strength across the periodic table, it fails to uphold when we compare HCl and HF or H₂S and H₂O. HCl is a strong acid and fully dissociates, but HF is a weak acid and only partially dissociates. Likewise, the acidity of H₂S is four times greater than the dissociation of H₂O. To explain this observation, we can use two theories: polarizability and ion-dipole solvation. HCl dissociates to form chloride and HF dissociates to form fluoride. In looking at the structure, we observe that chloride is larger than fluoride, which indicates that it has a smaller effective nuclear charge than fluoride and is thereby more polarizable—the electrons can distort to minimize electron-electron repulsions more readily in chloride than fluoride. This concept of polarizability is reiterated when comparing HCl (pK_a = -7), HBr (pK_a = -9), and HI (pK_a = -11). Iodide is the largest of the atoms in the set, therefore, it is more polarizable and has less electron-electron repulsions than observed in bromide or chloride making it the weakest conjugate base and HI the strongest acid.

The second consideration for explaining the trend in the hydrohalic series, stems from the strength of the ion-dipole interactions that result between water and the halides. For ion-dipole networks, the strength is inversely proportional to the internuclear distance between the two atoms of the bond. Therefore, iodide would form the weakest ion-dipole interactions and fluoride would form the strongest. In terms of enthalpy, the enthalpy change for dissociation is almost consistent amongst the four halides—the ion-dipole and polarizability effects are combined to rationalize why the change is almost consistent. Atoms with strong ion-dipole interactions are not as polarizable and vice versa. While enthalpy is consistent, the entropy effect is not consistent. Fluoride has very strong ion-dipole interactions and will form very rigid well-defined networks. Conversely, iodide, which has weaker ion-dipole interactions, will form less rigid more disordered intermolecular networks.

Fluoride (more rigid ion-dipole networks) Iodide (less rigid ion-dipole networks)

Figure 6: A Comparison the Ion-Dipole Networks Created by Fluoride and Iodide

We can rationalize that the dissociation of HI is more entropically favored than the dissociation of HF given that iodide ion-dipole networks are more disordered than fluoride ion-dipole networks.

To summarize, for binary acids, we use electronegativity to explain trends across the periodic table and size (ion-dipole interactions and polarizability) to explain acid/base trends down a column.
To rationalize the trends observed for other acids, we will need to use resonance and inductive effects. Let’s first consider the oxoacids, which have the form HXO_n that have hydrogen, oxygen, and a third atom such as sulfur, nitrogen, phosphorus, carbon, etc.

Once again, we can rationalize the strength by analyzing the stability of the conjugate base in each case. Let’s consider the conjugate bases of nitric acid (a strong acid that fully dissociates) and nitrous acid (weak acid). Nitrate the conjugate base of nitric acid has charge dissipated over three resonance forms and the nitrogen has a +5 oxidation state. Nitrite, which is a stronger base, only has two resonance forms and a +3 oxidation state. The greater number of resonance forms in nitrate indicate that charge is more dissipated (delocalized) on multiple atoms which increasing the stability and decreases the reactivity. Furthermore, the oxidation state can be used to rationale the attraction between the nitrogen atom and the charge on the respective oxygen atom. The nitrate structure has a +5 oxidation state indicating that it will attract charge more strongly than the +3 charge in nitrite, thereby, yielding greater stability from an electrostatic type interaction.

The final consideration for this discussion of structure involves inductive effects, which is the dissipation of electron density through sigma bonds because of the differences in electronegativity. This can be observed with oxoacids or other non-binary acids. Let’s consider the following examples:

The trifluoroacetic acid molecule has the smallest pK_a value indicating that it is the most acidic compared to acetic acid that is the least acidic. These structures all differ by the neighboring groups attached to the carboxylic acid structure. In each case, the halogen groups will withdraw electron density from the carboxylic acid through the carbon-carbon sigma bonds. If we analyze the conjugate bases of trifluoroacetate and acetate we observe both structures have the same number of resonance forms, the charges are located on the same atoms (oxygen) in the resonance forms, but the neighbors are differ. Trifluoroacetate is a very weak base (almost neutral). To understand the neighboring effect, we can analyze the electrostatic potential maps (Figure 7) of the conjugate bases. Recall that red regions indicate a large quantity of electron density, yellow regions indicate intermediate electron density, and blue regions indicate electro deficient regions.

Acetate Trifluoroacetate

Figure 7: The electrostatic potential maps for acetate and trifluoroacetate.

For acetate the oxygen atoms both have negative charges (through resonance) and the charges are localized on these two atoms. Conversely, with trifluoroacetate, we observe charges on the oxygens, but these are less significant than those observe in the acetate molecule. Furthermore, the green/yellow regions indicate the presence of charge illustrating that charge has been dissipated throughout the structure.

These effects are illustrated with hypochlorous acid, hypobromous acid, and hypoiodous acid.

pK_a

Cl-OH 7.5

Br-OH 8.7

I-OH 10.7

One final point to consider for acid-base strength is hybridization. Methylamine is sp³ hybridized (25% s-character) and is four times more basic than pyridine that is sp2 hybridized (33% s-character).

This is an important observation that will be reiterated at a later point. As the s-character increases, basicity decreases because s-orbitals lie closer to the nucleus than p-orbital; therefore, the charge or lone pair is held more tightly making it less reactive.

Note

To summarize: sp < sp² < sp³ In terms of base strength

sp > sp² > sp³ In Terms of acid strength

The molecular structure discussion focused upon ways in which the base electron density would be stabilized. The five factors included electronegativity, polarizability (solvation), inductive effects, resonance, and hybridization. To summarize:

For binary acids (H-A)

< >Electronegativity can be used to account for the increase in acid strength (decrease in base strength) across the periodic table. The greater the electronegativity, the stronger the attraction between the electrons in the base, and the more stable the resulting compound.Polarizability (Solvation) can be used to account for the increase in acid strength down a column in the periodic table (HI > HBr > HCl > HF in terms of acid strength). As the size of the atoms increases, the effective nuclear charge decreases, and the more polarizable the atom will become. This polarizability results in a reorganization of electron density to remove electron-electron repulsion and promote greater stability. This idea provides one account for the observation; however, we should also include entropy in the rationalization. With iodide, we will form weaker ion-dipole networks, but have more stabilization through polarizability, and we will have a greater (more positive) change in entropy because the ion-dipole networks will be less structured and less rigid. Fluoride forms very strong ion-dipole networks and dissociation is not entropically favored. Entropy dominates when explaining the acid strength of the hydrohalic acids.For Oxoacids (HXO_n or polyatomic acids, H-A-B)

< >Resonance/Oxidation Number. As the number of oxygen atoms is increased in an oxoacid, the oxidation number increases, and the number of resonance forms increases. This increase in oxidation number accounts for an increase in the attraction between the charge of the base and the parent molecule. Furthermore, the number of resonance structures is also increased which accounts for multiple ways in which charge is distributed. The charge is delocalized over multiple atoms that creates greater stability.Inductive Effects are used to explain variations in acid strength when the number of resonance forms and the oxidation numbers are exactly the same. Inductive effects describe the pull of electron density through sigma bonds by neighboring acids.

Example 7

Identify the strongest acid in each set and explain.

Set 1: H₂Se vs. AsH₃

Set 2: H₂Se vs. H₂S

Set 3: Cl-Se-H vs. H₂Se

Solution:

Set 1: H₂Se is a stronger acid because Se is more electronegative. Se will stabilize the resulting charge more strongly than As; therefore, H₂Se will be the stronger acid. Equilibrium favors the production of a weaker base.

Set 2: H₂Se is a stronger acid because Se is larger than S and is more polarizable and will create weaker ion-dipole networks which is entropically favored. Because of the polarizability and entropy effect, H₂Se will dissociate more than H₂S making it a stronger acid.

Set 3: Cl-Se-H is a stronger acid because of inductive effects. The Cl group will withdraw electrons through the sigma bonds that will dissipate electron density and make the conjugate base more stable. Equilibrium favors the production of a weaker base; therefore, Cl-Se-H will dissociate more readily than H₂Se.

Solvent Leveling Effect in Water

The strong hydrohalic acids, HI, HBr, and HCl, at the same concentration, cannot be differentiated in water.

HI (g) + H₂O (l) à H₃O⁺ (aq) + I^- (aq)

HBr (g) + H₂O (l) à H₃O⁺ (aq) + Br^- (aq)

HCl (g) + H₂O (l) à H₃O⁺ (aq) + Cl^- (aq)

The three acids, at 1 M concentrations, would all produce 1 M H₃O⁺ and would have a pH of 0. This observation would apply to the other monoprotic strong acids as well. This observation is called solvent – leveling which notes that the strongest acid that can exist in water is hydronium. All strong acids are levels to hydronium. The pK_a values listed in the table are estimates made in other solvents--dimethyl sulfoxide is a common solvent for these measurements.

The same observation can be made for bases when added to water. They automatically dissociate to produce hydroxide. Acids that are weaker than water (pK_a > 14) do exist, and they do form conjugate bases stronger than hydroxide, but once we add these to water, they revert to hydroxide. When solvent leveling is occurring, thermodynamics favors product formation (K > 1, DG < 0).

Given the solvent leveling effect of water, this chemistry does readily occur, but we may need to alter our solvent to prevent leveling to hydroxide or hydronium. For example, if we shift to ammonia as the solvent, bases stronger than hydroxide can be observe. Recall auto ionization occurs for all solvents, the strength of the acid and base resulting from auto ionization will determine the maximum strength of acids and bases that can be observed in those solutions.

In ammonia, the strongest acids that can ever be observed is ammonium (NH₄⁺) and the strong base is amide (NH₂^-). Acids stronger than ammonium will be leveled to ammonium. Bases stronger than amide will be leveled to amide.