There are three definitions for an acid that are commonly used: an Arrhenius acid, a Brønsted-Lowry acid, and a Lewis acid. Arrhenius acids are any substances that create hydrogen ions in solution, and Lewis acids are any substances that accept an electron pair. A Brønsted-Lowry acid is defined as any substance that donates a proton.

\(\displaystyle LiAlH_4\) is not an acid when it is used as a reducing agent. This can confuse students because, \(\displaystyle LiAlH_4\) is in fact donating a hydrogen ion to a compound that is being reduced; however, we are reducing something, and thus we are adding electrons. The only way that \(\displaystyle LiAlH_4\) could be an effective reducing agent is if it carried electrons on its donated hydrogen.

\(\displaystyle LiAlH_4\leftrightharpoons Li^++AlH_4^-\)

\(\displaystyle AlH_4^-\leftrightharpoons Al^{3+}+4H^-\)

As a result, we are donating hydride (\(\displaystyle H^-\)) ions, and not protons. Acids are exclusively those species that donate protons in solution, not hydrides.

The definition of a Lewis acid is an electron acceptor. Try drawing the Lewis dot structures for these compounds, and you will find that NH₄⁺, Na⁺, and Al³⁺ are missing electrons. They each have an empty orbital, allowing them to accept electrons.

BF₃ is not an ion; however, we know that boron only has 3 valence electrons, which means that even when they are all bound to fluorine the molecule does not satisfy the octet rule. BF₃ only has 6 valence electrons around boron, and can accept another electron pair to get to the octet state.

CaO is a Lewis base. In the Lewis dot structure, we can see that the oxygen molecule has two lone pairs that it can donate to other molecules, making it an electron donor.

First, we need to determine what type of reaction HCO₃^- is participating in. We can see that a H⁺ is being produced in the forward reaction, and consumed in the reserve reaction; thus, we are looking at an acid-base reaction. H₂CO₃ has lost a H⁺, making it a Arrhenius acid. HCO₃^- is the conjugate base of the acid H₂CO₃.

The stronger the acid, the weaker its conjugate base. This is because a strong acid will dissociate almost entirely, meaning that its conjugate base will not accept protons frequently. The only strong acid as an option is nitric acid (HNO₃). As a result, we conclude that nitric acid has the weakest conjugate base.

Certain salts are able to affect the pH of a solution, because the dissociated ions can act as acids or bases. The conjugate bases of strong acids will not affect the pH, nor will the conjugate acids of strong bases.

NH₄⁺ is the conjugate acid of ammonia (NH₃), a weak base. As a result, it will donate protons to the solution and make the solution acidic.

A diprotic acid will ionize fewer of its second protons than its first; that is, a smaller fraction of the \(\displaystyle HSO_{4}^{-}\) ions will react to \(\displaystyle SO_{4}^{2-}\) than \(\displaystyle H_{2}SO_{4}\) molecules will react to \(\displaystyle HSO_{4}^{-}\).

\(\displaystyle H_2SO_4\rightarrow HSO_4^-+H^+\)

\(\displaystyle HSO_4^-\rightarrow SO_4^{2-}+H^+\)

The presence of protons in solution from the original ionization of \(\displaystyle H_{2}SO_{4}\) in the first reaction drives the equilibrium toward \(\displaystyle HSO_{4}^{-}\) (reactants) and away from \(\displaystyle SO_{4}^{2-}\) (products) in the second reaction.

The question is asking which of the following is a strong acid. Of the following, the only strong acid provided is perchloric acid, HClO₄. NH₃ is a weak base; HCOOH, C₆H₅COOH, and HCN are all weak acids. A strong acid is one that dissociates completely in water.

The strong acids that you must know for the MCAT are:

Hydroiodic acid (HI)

Hydrobromic acid (HBr)

Hydrochloric acid (HCl)

Perchloric acid (HClO₄)

Sulfuric acid (H₂SO₄)

and Nitric acid (HNO₃)

Hydroiodic acid (HI) is the strongest acid listed. Charge density decreases as the atomic size gets larger, thus stabilizing the charge when a hydrogen is given off. Hydroflouric acid (HF) is a relatively weak acid because electronegative elements hold on to their valence electrons more tightly, and are thus less likely to dissociate.

Strong bases are those that completely dissociate in water. All group I hydroxides are considered strong bases, including sodium hydroxide, potassium hydroxide, and lithium hydroxide. Some group II hydroxides are also strong bases, including calcium hydroxide, barium hydroxide, and strontium hydroxide.

Ammonia is considered a weak base and is frequently used as a buffer in solution with ammonium, its conjugate acid.