Acids and Bases

Acids and bases are common classes of chemicals that react with each other and with water. As a result, they are important biologically, industrially, and environmentally.

Bases are bitter and feel slippery to the touch. Common bases are found glass cleaners (ammonia), antacid tablets (carbonate), baking soda (also carbonate), and toothpaste (fluoride).

Acids are often sour tasting. Acids also react with some metals, releasing \(\ce{H2}\) gas and causing corrosion. Common acid-containing materials include oranges (citric acid), vinegar (acetic acid), aspirin (acetylsalicylic acid), car batteries (sulfuric acid), and acid rain (sulfuric acid and nitric acid). Acid rain damages buildings and artwork in cities around the world.

Many biological molecules are also acids or bases, though their structures may be more complex and difficult to identify at first. The purines and pyrimidines in DNA are bases, while dietary fats are acids.

The strength of an acid or base depends on its ability to ionize. Strong acids and bases dissociate completely. None of the original molecule remains when it reaches equilibrium in an aqueous solution. Weak acids and bases reach an equilibrium where the majority of the molecules remain in their original form.

Dissociation of nitric acid, a strong acid.

\[\ce{HNO3 (aq) -> H+ (aq) +NO3- (aq) }\]

Dissociation of hydrofluoric acid, a weak acid. \[\ce{HF (aq) <-> H+ (aq) + F- (aq) }\]

The first modern definition of acids and bases was written at the end of the 1800's. Under the Arrhenius definition, acids produce a proton, \(\ce{H^+}\) , in solution, while a base is any substance that produces a hydroxide ion, \(\ce{OH^-}\). When an acid and base react with each other, the products are always a salt and water.

The dissociations and reaction of HCl (a strong acid) and NaOH (a strong base) is illustrated below.

\[\ce{HCl (aq) -> H+ (aq) + Cl- (aq) }\] \[\ce{NaOH (aq) -> Na+ (aq) + OH- (aq)}\] \[\ce{HCl (aq) + NaOH (aq) -> H2O (l) + NaCl (aq)}\]

While this simple definition of acids is useful, it is also limited in that it only applies to reactions that take place in water and to compounds that produce a proton or a hydroxide ion. Acidic or basic behavior can be observed in many systems that do not fit those requirements.

In the 1920's, a more general definition of acids was proposed. Brønsted–Lowry acids are species that donate a proton and bases are species that accept a proton. Brønsted–Lowry acids and bases must always be identified as a conjugate pair, because the electrons are moving from one species to another.

Consider the following equation:

\[\ce{NH3 (g) + HCl (g) -> NH4Cl (s) }\]

Would this be an acid-base reaction under the Arrhenius definition?

No. \(\ce{HCl}\) is an acid, but there is no base to form \(\ce{OH^-}\).

Would this be an acid-base reaction under the Brønsted–Lowry definition?

Yes. There are two conjugate acid-base pairs: \[\ce{NH3 (g) -> NH4+}\] Ammonia, \(\ce{NH3}\), is a base because it is receiving a proton to form its conjugate acid, the ammonium ion \(\ce{NH4+}\).

The other reactant, \(\ce{HCl}\), is both an Arrhenius acid and a Brønsted–Lowry acid. Its conjugate base is \(\ce{Cl-}\). \[\ce{HCl -> Cl-}\]

The Lewis definition of acids and bases was published around the same time as the Brønsted–Lowry definition. It is very similar to the Brønsted–Lowry definition, except that The Lewis definition hinges on the movement of electrons rather than protons. A Lewis acid accepts an electron pair. A Lewis base donates an electron pair. The Lewis definition is the most inclusive definition of acids and bases.

Are there any Brønsted–Lowry acids that are not also Lewis acids?

No. Every Brønsted–Lowry acid (a species that donates \(\ce{H^+}\)) also meets the definition of a Lewis acid (a species that accepts electrons). Similarly, all Brønsted–Lowry bases are Lewis bases.

Are there any Lewis acids that are not Brønsted–Lowry acids?

Yes. \(\ce{AlCl3}\) and \(\ce{BCl3}\) are two examples of species that can accept electrons, but do not have protons (\(\ce{H^+}\)) to donate during a reaction.

The Usanovich theory is one of the least used due to several factors and therefore little known but very important to explain acid - base reactions involving oxirreduçao.Nesta theory the Russian chemist claims that the acid and any compound that can donate electrons , receive anions and donate protons and every base is any compound that can receive electrons receive cations and receive protons.

\[\ce{2 MnO4^- + 5 H2O2 + 6 H^+ -> 2 Mn^2+ + 8 H2O + 5 O2 }\]

While it is common to see the product of an acid dissociation written as \(\ce{H+}\), protons are too reactive to remain alone in solution. They combine with water to form the hydronium ion. The dissociation of \(\ce{HCl}\) would more accurately be written as follows.

\[\ce{HCl (aq) + 2 H2O (l) -> H3O+ (aq) + Cl- (aq) + OH- (aq)}\]

True or false: the above equation is equivalent to writing \[\ce{HCl (aq) -> H+ (aq) + Cl- (aq) }\]

True. It seems that we have picked up an extra reactant (water) and product (hydroxide), but as we'll see momentarily, these species can be neglected in the overall balanced reaction.

True or false: in a beaker of distilled water, all the water molecules take the form \(\ce{H2O}\)

False. Pure water does not only contain \(\ce{H2O}\). Water is constantly undergoing chemical reactions where one molecule of water is acting as an acid and another is acting as a base, called autoprotolysis.

\[\ce{ 2 H2O (l) <-> H3O+ (aq) + OH- (aq)}\]

The concentration of \(\ce{H3O+}\) in pure water is approximately \(\ce{1*10^{-7}}\) moles per liter at room temperature, as is the concentration of \(\ce{OH-}\).

Since autoprotolysis is a constant process in any aqueous solution, it can be disregarded (or considered as a separate reaction) in simple acid-base problems.

The concentration of hydrogen ions is closely monitored in soil, water, industrial processes, various types of waste, and the human body, among other places. Hydrogen ions concentrations in solutions are often very small, but can quickly change at an exponential rate. A logarithmic scale is therefore useful in quantifying these changes.

Hydrogen ion concentration is measured as pH, which is defined as:

\[\ce{pH = - log [H^{+}] = log \frac {1}{[H^{+}]}}\]

Similarly, hydroxide can be measured on a pOH scale. \[\ce{pOH = - log [OH^{-}] = log \frac {1}{[OH^{-}]}}\]

Autoprotolysis, also called autoionization, involves the dissociation of water to hydrogen and hydroxide. Writing that equation in its simplest form results in the following: \[\ce{H2O (l) -> H+ (aq) + OH- (aq)}\]

This reaction is a chemical equilibrium with the dissociation constant Kw.

\[\ce{Kw = [H+][OH^{-}] = 10^{-14}}\]

Rewritten in logarithmic form:

\[\ce{pH + pOH = 14}\]

What is the pH of pure water?

For every one mole of water that dissociates, one mole of hydrogen ions is generated, and one mole of hydroxide ions is generated.
\[\ce{H2O (l) -> H+ (aq) + OH- (aq)}\]

\[\ce{pH + pOH = 14}\] \[\ce{pH = pOH}\] \[\ce{pH = 7}\]

pH indicators are dyes that change color at different concentrations of hydronium. Liquid and paper versions are available, and a cabbage-based pH indicator can be made at home. Some change color once at a specific pH (indicating when a solution has moved from being acidic to basic, for example) while others can show a spectrum of colors and measure pH from 1 to 14. Top image [2], bottom image [3].

One place where the pH scale is useful is in environmental science. Rivers contain water with a pH between 6 and 9. If the pH of a river changes, the person measuring it knows not only that a contaminant exists, but also has some idea of what it is. Most animals and their eggs or larva can only survive within a certain pH range. If the river's pH changes, ecologists can make good predictions regarding which animals will be negatively affected, and which new species might move in to the area to take their place.

[1] Image from https://commons.wikimedia.org/wiki/File:Pollution-Damagedbyacid_rain.jpg under Creative Commons licensing for reuse and modification.

[2] Image from https://commons.wikimedia.org/wiki/File:PHindicatorpaper_roll.jpg under Creative Commons licensing for reuse and modification.

[3] Image from https://commons.wikimedia.org/wiki/File:ColoredpHindicator_(circle).png under Creative Commons licensing for reuse and modification.