The reaction between acids and bases depends on the role protons play and what the molecules react with. For example, we know that if a molecule has a proton that it really wants to get rid of, this molecule is considered an acid. On the other hand, if we have another molecule that only somewhat readily accepts a proton, then it is considered a weak base.
Understanding these trends and how acids and bases react with each other is critical to succeeding in AP Chemistry. So, let's dive into the molecular structures of acids and bases!
Molecular Structure, Bonding and Acid-base Behavior
Before diving into the molecular structures of acids and bases, let's review what acids and bases are. Depending on which chemist you ask, acids and bases have different definitions!
According to Swedish chemist Svante Arrhenius, acids and bases are defined based on their ability to donate H+ ions or OH- ions in water (H2O).
Arrhenius acids are H+ donors in water (H2O).
Arrhenius bases are OH- donors in water (H2O).
Another definition of acids and bases is given by Johannes Brønsted and Thomas Lowry. According to them, acids and bases differ in their ability to donate or accept H+ ions.
Brønsted-Lowry acids are H+ donors.
Brønsted-Lowry bases are H+ acceptors.
Finally, Gilbert Lewis, an American chemist, came up with a definition of acids and bases based on electron donation or acceptance.
Lewis acids are electron acceptors.
Lewis bases are electron donors.
Acid-base behavior (whether a molecule behaves like an acid or a base), and its strength is greatly affected by its molecular structure and bonding. Let's start by talking about the molecular structure of acids and bases.
Molecular Structure and Acid-Base Strength
There are two different types of acids you need to know about: binary acids, and oxyacids. Let's start with binary acids (figure 1). Binary acids are acids with the formula H—X, where a hydrogen atom is bonded to an electronegative nonmetal atom, X.
- X is typically a halogen (group 7 element), but can also be nitrogen (N), phosphorus (P), sulfur (S), selenium (Se), and tellurium (Te).
Figure 1. Binary Acid Examples, Isadora Santos - StudySmarter Originals.
Next, we have oxyacids. Oxyacids are acids that have one (or multiple) H-O bonds, where a hydrogen atom is bonded to an oxygen atom. Nitric acid (HNO3) and sulfuric acid (H2SO4) are oxyacids. The common molecular structure of an oxyacid is \( \text{H}^{+}\text{ + }\text{nonmetal + oxygen (O)} \).
Effect of Structure on the Strength of Acids and Bases
Binary acid strength can depend on electronegativity or atom size (figure 2). Typically, when it comes to binary acids, the stronger the bonding between H and X, the weaker the acid.
↳ HI > HBr > HCl > HF
For elements in the same period, we look at electronegativity: the more electronegative the atom, the stronger the acid. This is because the more electronegative an atom is, the more it likes retaining its electrons!
↳ HF > H2O > NH3
Figure 2. Binary acid strength trend, Isadora Santos - StudySmarter Originals.
Let's look at an example!
Which of the following compounds have the greatest binary acid strength?
$$ NH_{3}\text{ vs. }CH_{4} \text{ vs. }HF $$
The first thing you need to do is look at the atoms bonded to H in each compound. In this case, N, C, and F are in the same period. So, that means we will use the trend in electronegativity to find out which compound is the most acidic.
The most electronegative atom (F) will be the most acidic. Therefore, HF is the compound with the greatest binary acid strength.
The higher a bond's enthalpy is, the more energy is required to break it. A high bond enthalpy means a strong bond, and therefore, a less acidic compound. To learn more about this, check out "Enthalpy"!
In terms of strength, the strength of oxyacids is based on the number of oxygen atoms present or the electronegativity of the nonmetal (figure 3).
When dealing with compounds having a different number of oxygen atoms, we say that the more oxygen atoms, the more acidic.
When dealing with compounds that have the same number of oxygen atoms, then the more electronegative the nonmetal, the more acidic.
Figure 3. Examples of Oxyacids, Isadora Santos - StudySmarter Originals.
You are asked to explain to your peer why H2SO4 is a stronger acid than H2SO3. How would you go about explaining this?
This can be answered through the oxyacid extension of the electronegativity rule. The additional oxygen atom in H2SO4 provides additional electronegativity, and therefore, a greater electron density, and a more stable conjugate base. The more electronegative the group attached to the hydrogen atom is, the more acidic the molecule is overall. Or, in other words, the more stable the conjugate base is, the stronger the acid will be. H2SO4's conjugate base being stable is consistent with it being a stronger acid.
Did you know that molecular structure also affects the pH of acids and bases? Basically, when acidic compounds lose a proton (H+ ion), the pH increases, whereas when a base gains a proton, the pH decreases!
Examples of Molecules with given Lewis structure Acid and Base
For example, let's look at the Lewis structure of sulfuric acid and nitric acid (figure 4). In acid-base reactions, acids tend to donate an H+ ion to the base.
Figure 4. The lewis structure of nitric acid and sulfuric acid, Isadora Santos - StudySmarter Originals.
Compared to acids, bases almost always have the structure of MOH where M is a metal cation. However, they are some exceptions such as NH3, which is a non-hydroxide base. It is thanks to its molecular structure that bases are able to accept a proton (H+).
Let's look at the Lewis structure of the base ammonia (NH3) in figure 5. To be able to accept a proton (donated from the acid) and form a bond, bases must have an unshared pair of electrons (shown as blue dots). For instance, in a chemical reaction between ammonia (NH3) and hydrochloric acid (HCl). the acid donates an H+ ion to ammonia. When this happens, the bond between H and Cl breaks.
Figure 5. Chemical reaction between ammonia and hydrochloric acid, Isadora Santos - StudySmarter Originals.
Common strong bases include sodium hydroxide (NaOH), potassium hydroxide (KOH), and other group 1 and 2 metals bonded to one or more hydroxides (based on the charge of the metal.) An example of a non-hydroxide base is ammonia (NH3).
If you understand the rules behind how acids structurally appear, understanding bases will be easy to comprehend! Remember, acids and bases perform complementary roles, meaning that the rules that dictate how acids structurally appear are just reversed for bases.
- The less electronegative a molecule is, the more basic it is.
- The weaker a conjugate acid, the more strong a base is.
- The stronger the bonds within a molecule, the more likely it is to be basic.
What we can gather from this is a simple rule: acids and bases are opposites in strength. This means that a strong acid would be a weak base, and a strong base would be a weak acid. As long as you can grasp one set of structural rules, the other is intuitive!
Atomic and Molecular Structure Equilibria Acids and Bases
Acid-base equilibrium reactions prefer the side of a reaction that has weaker acids and bases. This is because stronger acids and bases can perform their roles of donating and accepting protons more readily, making them much more likely to react than weaker acids and bases!
If these weaker acids and bases aren't reacting, this means they are staying in their weak forms, making them the preferred product of our equilibrium problem.
We can apply the structural rules that we just learned about acids and bases to determine which side of an equilibrium reaction will be preferred. Let's try a simple example.
Which side of this acid-base equilibrium reaction will be preferred?
$$ \text{HF + NH}_{3}\text{ } \rightleftharpoons \text{ }\text{NH}_{4}^{+} \text{ + F}^{-} $$
Recall that we want to find out which side of our reaction has the weaker acid. We can identify HF as the acid and NH3 as the base, meaning that NH4+ is our conjugate acid and F- is our conjugate base. Now, it's simply a matter of comparing acid strengths. Using our structural rules, we know that hydrofluoric acid will be much more strong than the conjugate acid, NH4+. This means that equilibrium will favor the right side of this reaction.
As you learn more about how acid-base equilibria work in the remainder of AP Chemistry and higher-level chemistry courses, these trends will reappear again and again. Therefore, if you understand how acid and base structure influences their functionality, you'll be able to comprehend not just the how, but behind acid-base equilibria!
Molecular Structures of Acids and Bases - Key takeaways
- Binary acids are acids with the formula H—X, where a hydrogen atom is bonded to an electronegative nonmetal atom, X.
- Binary acid strength can depend on electronegativity or atom size
- Oxyacids are acids that have one (or multiple) H-O bonds, where a hydrogen atom is bonded to an oxygen atom.
- The strength of oxyacids depends on the number of oxygen atoms present in the compound or the electronegativity of the non-metal atom.
- Bases almost always have the structure of MOH where M is a metal cation. However, they are some exceptions such as NH3, which is a non-hydroxide base.
References
- Zumdahl, S. S., Zumdahl, S. A., & Decoste, D. J. (2019). Chemistry. Cengage Learning Asia Pte Ltd.
- Theodore Lawrence Brown, Eugene, H., Bursten, B. E., Murphy, C. J., Woodward, P. M., Stoltzfus, M. W., & Lufaso, M. W. (2018). Chemistry : the central science (14th ed.). Pearson.
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