The Ka is the acid dissociation constant, and thus it is what determines how strong the acid is. Stronger acids dissociate to a greater extent and produce lower pH values.

All of the listed salts will dissolve into ions when in water. When the ions are in solution, they can act as acids or bases by donating or accepting protons. Chloride, bromide, and iodide ions are all conjugate bases of strong acids, so they will not accept protons. Sodium and potassium ions are the conjugate acids of strong bases, which dissociate completely, so they will not accept hydroxide ions.

Ammonium is the conjugate acid of ammonia, a weak base. The ammonium ion can donate a proton to the solution. This will make the solution slightly acidic. As a result, ammonium bromide is a salt that will make an acidic solution.

A Lewis acid is a 2-electron acceptor, and a Lewis base is a 2-electron donor.

, while it is acidic, cannot accept two electrons, so it is not a Lewis acid.

can accept two electrons, so it is a Lewis acid.

is a Lewis base.

, like hydrochloric acid, is acidic, but cannot accept two electrons, so it is not a Lewis acid.

A Lewis base is an electron-pair donor. Only PH3 has a pair of nonbonding electrons and can act as a donor.

For this question, we'll need to understand the different definitions of an acid in order to answer it. There are three definitions of acids that are important to know.

1. Arrhenius acids

These are compounds that, when added to water, increase the concentration of ions present in solution.

2. Bronsted-Lowry acids

These are any acid that can release , even while not in water.

3. Lewis acids

This is the most general definition of acids. It is any compound that can accept a lone electron pair.

Lewis acids are the most general kind of acids, meaning that any acid that is Bronsted-Lowry or Arrhenius will also be a Lewis acid. However, the reverse is not true. Not all Lewis acids will fall under the category of Bronsted-Lowry or Arrhenius.

The correct answer in this question is aluminum chloride. We can see that, based on aluminum's position in the periodic table, it has three valence electrons in its outer shell. Each of these electrons is tied up in a shared bond with a chloride. This means that the aluminum in aluminum chloride has six valence electrons. However, since aluminum has a maximum capacity of eight valence electrons, it has room for two more. This vacancy allows the aluminum component of aluminum chloride to accept an electron pair from any sort of electron donor. Thus, aluminum chloride qualifies as a Lewis acid. However, aluminum chloride has no way of producing . Consequently, it is neither a Bronsted-Lowry acid nor is it an Arrhenius acid.

Weaker acids dissociate less in water and therefore, reverse reaction is favored in

. This indicates that the weaker the acid, the stronger its conjugate base.

, and are all strong acids. Their conjugates are weak bases.

has the weakest conjugate acid among the all the choices, so it is the strongest base.

The Bronsted-Lowry definition of an acid is a substance that can donate a hydrogen ion and forms its conjugate base; a Bronsted-Lowry base accepts a hydrogen ion and forms its conjugate acid. Thus we are looking for a substance that can either donate or accept a hydrogen ion (amphoteric). Bisulfite may give up a proton to become , a Bronsted-Lowry base. It acts as an acid as , which can donate up to two hydrogens.

Bases can be defined as species that quench hydrogen ions from a solution. A hydroxide ion and a hydrogen ion combine to form water in solution. Recall that basis solutions range in pH from about 7 to 14.

Every Lewis acid is also a Brønsted-Lowry acid, and every Brønsted-Lowry acid is an Arrhenius acid; thus, H2SO4 is all three, since it is an Arrhenius acid: (it dissolves in water to produce a proton). Sulfuric acid is also considered a strong acid, as it full dissociates in water.

The PRODUCTS of a neutralization reactions are salt and water, not the reactants. The rest of the options all correctly pertain to neutralization reactions.