pH and POH of Strong Acids and Bases - AP Chemistry

The definition of a Bronsted-Lowry base is a species that has the ability to gain, or accept a proton (H+). Dissociating in solution is part of the Arrhenius definition of acids and bases, and Lewis acid are electron pair acceptors.

A Lewis acid is a species that can accept an electron pair. In the BCl3 molecule, B does not have a complete octet (3 covalent bonds, thus 6 electrons around it rather than 8). Thus, it can accept another electron pair, making it a Lewis acid.

Conjugate acid has one more H+ than the compound with which it is being compared. Thus, H2CO3 is the conjugate acid of HCO3–.

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.

HCN, HCl, CH3COOH, and NH4Cl are all acids (NH4+ is the ammonium ion). That only leaves KCN as the correct answer.

Both are strong acids, but H2SO4 is bivalent, realeasing 2 protons for each molecule dissolved in solution. Further, a more acidic solution would have a lower pKa.

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 conjugate base of an acid is the same formula, minus one proton. The only option that fits this description is the one that includes the hydronium ion (H3O+) and water (H2O).

The hydroxide ion (OH–) is a strong base and therefore would not be present in an acidic solution. Protons, the hydronium ion, and water will all be present in relatively large amounts within the solution.

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

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.

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.

The above reaction describes the reaction that occurs when carbon dioxide dissolves in water and reacts to form carbonic acid, which is responsible for acid rain.

The phenomenon seen by dry ice is simply the sublimation of carbon dioxide. Acids and bases do not usually react in a violent fashion. The reaction between baking soda (sodium bicarbonate) and vinegar (acetic acid) is given below. The production of carbon dioxide gas is responsible for the bubbles seen in this reaction.

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.

In order to find the base dissociation constant for the conjugate base, we can start by finding the acid dissociation constant for the acid. Since a 1M solution of the acid has a pH of 4.6, we can find the proton concentration of the solution.

Since the acid is monoprotic, we can set the following equilibrium expression equal to its acid dissociation constant.

We can see that, since the acid is monoprotic, the concntration of protons will be equal to the concentration of the acid anion. The final concentration of the acid molecule will be equal to the initial concentration, minus the amount of protons formed. Using these values, we can solve for the equilibrium constant for the acid.

Now that we have the acid dissociation constant, we can find the conjugate base's dissociation constant by setting the product of the two values equal to the autoionization of water.

A salt will dissociate completely in an aqueous solution, forming its respective ions. Knowing this, we can make predictions on how the ions will affect the pH of the solution, given their strengths as conjugate acids and bases. Fluoride ions are the conjugate base of hydrofluoric acid, which is a weak acid. As a result, fluoride ions will attach to protons in solution and decrease the proton concentration of the solution. This will make the solution more basic.

Chloride and bromide ions are conjugate bases of strong acids. Since these strong acids would dissociate completely in water, these ions will not associate with hydrogen ions in solution and will not affect the solution pH.

Acids can be defined as species that donate hydrogen ions to solutions. If there is a hydrogen group on a molecule, it is possible that it may be donated to the solution, which will result in a decrease in pH.

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.

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.