Chapter 15 Acids and Bases

Brønsted Acids and Bases

A Brønsted acid is defined as a substance that is capable of donating a proton, and a Brønsted base as a substance that is capable of accepting a proton.
These definitions lead to the conjugate acid-base pair concept. The conjugate base in this instance is the substance that remains when one proton has been taken away from a Brønsted acid. The opposite can be said for a Brønsted base and its conjugate acid. This is demonstrated in the example reaction below:


In this reaction, NH3 serves as a base (symbolized by a b) because it accepts a proton (H+ ion) as the reaction progresses and becomes NH4+, its conjugate acid (symbolized by an a). The same goes for HCl, except it is classified as an acid because it loses a proton, and becomes Cl-, its conjugate base (the pair is also classified with an a and b and conjugate pairs linked with arrows).

The Acid-Base Properties of Water

Water has a special property which is its ability to act either as an acid or as a base. It functions as a base in reactions with acids, and as an acid in reactions with bases.

The Ion Product of Water

The pH Scale


Strength of Acids and Bases


Conjugate Acid-Base Pairs

There are certain properties which exist among conjugate acid-base pairs, which are the following:
From Chemistry: Eighth Edition, Raymond Chang

Weak Acids and Acid Ionization Constants

Example from Chemistry: Eighth Edition, Raymond Chang

Weak Bases and Base Ionization Constants


The Relationship Between the Ionization Constants of Acids and Their Conjugate Bases


Diprotic and Polyprotic Acids

The ionizations of polyprotic acids are treated in a similar fashion to that of monoprotic acids; however, they ionize in steps. Basically, they lose one proton at a time, and an ionization constant must be written for each step. A table of some polyprotic acids and their breakdown steps is shown below.

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(Table from Cliffsnotes Chemistry)


Molecular Structure and the Strength of Acids

An acid’s strength depends on several factors, such as temperature, solvent properties, and its molecular structure. Considering a certain simple acid HX, we measure its strength by its tendency to ionize. There are two factors that influence to what extent this acid will ionize. They are the strength of the H—X bond, and the polarity of the H—X bond. A stronger bond means that it is more difficult to break, and therefore the acid will not ionize to a great extent, and therefore be weak. If the bond is very polar, then there will be a large gathering of positive and negative charges on the H and X atoms, allowing it to split into both positively charged H ions and negatively charged X ions. Therefore, a high polarity results in a stronger acid. Now let us consider several specific types of molecular structures of acids.

Hydrohalic Acids
The halogens form a series of monoprotic acids known as hydrohalic acids (HF, HCl, HBr, HI). If we examine the bond energies of each of these species, it would be noticed that HF has the highest, and it decreases as we move down the group towards I. Therefore when it comes to hydrohalic acids, HF is the weakest (due to its bond energy), and they increase in strength up to HI. So from weakest to strongest, these acids are HF, HCl, HBr, and HI.

Oxoacids contain hydrogen, oxygen, and one other element, which we will call X. The basic structure of these molecules is characterized by the presence of one or more hydrogen (O—H) bonds. Now, X may either be a single atom, or a polyatomic structure which is attached to the hydrogen bonded oxygen. Due to the presence of this hydrogen bond, bond strength provides negligible influence on acidity. Therefore, the electronegativity of the atom X is what will determine the acidity. If X is an element of high electronegativity, it will attract more electrons, therefore making the X—O bond more covalent. This results in a decrease in strength of the hydrogen bond, allowing the molecule to dissociate into negatively charged XO ions and positively charged hydrogen ions.

To even further examine these types of acids, we can divide them into two distinct groups:

1. Oxoacids with different central atoms, X, that are from the same periodic group (same oxidation number). For example, the oxoacids in which X is either N, P, As, or Sb all fall into one group. In this group, the strength of the oxoacid increases as we move up a periodic group (like the group of atoms mentioned before the oxoacid with Sb would be strongest, and the one with N would be weakest).
2. Oxoacids with the same central atom, X, but different numbers of attached groups. This includes acids like hypochlorous acid, chlorous acid, chloric acid, and perchloric acid. They all have chlorine as the central atom, but have different numbers of other atoms (oxygens in this case) attached to it. The electronegativity of the central atom increases with the number of attached atoms, so the more attached atoms X has, the stronger the acid will be.

Carboxylic Acids

Carboxylic Structures are organic acids with Lewis structures represented as below:
external image 748px-Carboxylic-acid.svg.png

(Image from Wikipedia)
In this structure, R is part of the acid molecule and the other parts represent the carboxyl group (COOH). The strength of these acids depends on the acidic nature of R. Such a difference can be explained below with both acetic and chloroacetic acids:
(Images from Google image searches)

Notice that chloroacetic acid is stronger than acetic acid, even though they differ only by the presence of one chlorine atom replacing a hydrogen atom. The reasoning for this is simple; the presence of this extra chlorine causes the electron density to shift towards the R group, causing the hydrogen bond (O—H) to become more polar. This allows the chloroacetic acid molecule to ionize to a greater extent than acetic acid, therefore making it stronger by definition.

Acid-Base Properties of Salts

By definition, we know that a salt is an ionic compound formed from the reaction between an acid and a base. Salt hydrolysis is the term used to describe the reaction of an anion or cation of a salt with water. This usually affects the pH of a solution. Salts that contain a cation which is either an alkali metal or an alkaline earth metal and an anion which is the conjugate base of a strong acid (ex: NaCl and KBr), produce a neutral solution. Ions from these two groups both have little tendency to react with water, and are not hydrolyzed. Salts that produce basic solutions contain either an alkali or alkaline earth and the conjugate base of a weak acid. We can do another calculation when it comes to this type of solution: percent hydrolysis. To do this we need to use the following formula:
Finally, we can use this x value to obtain the pOH (which is 4.89), the pH (9.11, which is what we expect as this is a basic solution), and the percent hydrolysis, found by using the formula:

That is the general formula for dealing with acid-base calculations for salt solutions. Everything else that we would need to know is what kinds of salts make up acidic solutions, and salts in which both the cation and anion hydrolyze. For the first one, the cation must be the conjugate acid of a weak base, and the anion must be the conjugate base of a strong acid. In this case, it is the opposite of a base in terms of hydrolysis: the cation is hydrolyzed and the anion is not. However, if both the cation and anion are from weak bases, they will both hydrolyze. The pH of such a solution is not always 7, as is depends on the relative strength of the weak acid and base. A summary of all the possible examples can be found in the following table:

(From Chemistry: Eighth Edition, Raymond Chang)

Acid-Base Properties of Oxides and Hydroxides

An oxide compound can be classified as basic, acidic, or amphoteric. It should be noted that all alkali metal oxides and alkaline earth metal oxides (except BeO) are basic. BeO and several other metallic oxides from group 3A and 4A are actually amphoteric. Nonmetal oxides of which the oxidation number of the representative element (other element besides oxygen) is high are acidic, and of which the oxidation number of that element is low are basic. To get an overall view of which elements form acidic, basic, and amphoteric oxides, look at the table below:
There are a few types of reactions that these oxides can go through. The first is that basic metallic oxides and water react to form metal hydroxides, like in these examples:

The next is the reactions between acidic oxides and water, which form acid solutions, like in the following examples:

The last types of reactions are those between acidic oxides and bases, and basic oxides and acids. These will follow through like standard acid-base reactions, in that they produce a salt and water. The following is one example of this process:
Some oxides, however, are amphoteric. They can react as either an acid or a base, depending on what other compounds they react with. Here is an example using aluminum oxide:

Basic and Amphoteric Hydroxides
We know that the alkali and alkaline earth metal hydroxides have basic properties now. So now we can add the amphoteric hydroxides to this list. It’s also important to note that every one of these amphoteric hydroxides is insoluble. They include beryllium, aluminum, tin, lead, chromium, copper, zinc, and cadmium hydroxide. The following example shows that aluminum hydroxide can react with both acids and bases:


Lewis Acids and Bases

Up to now, we have been using only the Brønsted definition of acids and bases, which defines an acid as a substance that can donate protons, and a base as a substance that can accept protons. If we were to examine the Lewis structures of two well known bases, the hydroxide ion and ammonia, we would notice that the atom the extra proton attaches itself to has at least one pair of unshared electrons. This is a category shared amongst almost all other bases. So, in 1932, an American chemist, Gilbert N. Lewis proposed a new definition of acids and bases which was based on chemical bonding theory. His definition described an acid as a substance that can accept a pair of electrons (now known as a Lewis acid), and a base as a substance that can donate a pair of electrons (now known as a Lewis base). These are significant definitions in that they are much more general definitions of acids and bases, thus they can include more reactions than the Brønsted definitions. For example, we can consider the reaction between boron trifluoride and ammonia to form another compound:
It can be seen that the B atom in the boron trifluoride has no unpaired electrons. It therefore accepts the pair from the ammonia molecule, acting as an acid (and ammonia the base) according to the Lewis definitions. So basically, in order to determine which compound in an acid-base reaction is a Lewis acid or base, we need to examine their Lewis structures to see which compound is capable of accepting or donating a pair of electrons.


Calculating the pH and pOH of Acids and Bases
Purpose: To determine both the pH and pOH of various acids and bases using the ICF (initial, change, and final concentration) method.

Now, using the process shown above, fill out the data table below.

When finished, to check your answers, go to the ChemCollective virtual lab. Once there, note the side bar of solution selections. Find the species you wish to check under the solution tabs, then drag and drop it into the work space. Click once on the virtual beaker and some information will come up on the right side. This will tell you the concentrations of each species within the acid or base as well as the pH (from which you can discern the pOH). The whole setup of the virtual lab should be pretty self explanitory once you get there.

Review Questions

Questions from Chemistry: Eighth Edition, Raymond Chang

Answers to Questions

Resources Used
Chemistry: Eighth Edition, Raymond Chang

Cliffsnotes Chemistry

Google Image Search

ChemCollective Virtual Chemistry Lab