Introduction

Acids and bases




What is acids and bases?

When we talk about acids and bases in everyday life, we have a clear notion about what we mean. Or rather: We think we have a clear notion.

The most common definition for acids and bases for educational purposes it the one called Brønsted acids and bases (sometimes also spelled Bronsted). It is not the only definition, and most importantly, it does not cover all types of acids and bases, which means that you can very well have acidic or alkaline conditions, without the presence of an acid or base, if you stick with this definition.

Robert Boyle's definition of acids and bases
Before the concept of ions and the understanding of chemical compounds, which came later, the Englishman Robert Boyle defined acids and bases from functionalities.

Acids
Baser

Svante Arrhenius' definition of acids and bases
The Swede Arrhenius introduced the concept of ions in 1884, and in 1887 he enriched it to
This is the equivalent to the most common misunderstanding of what acids and bases are.


Johannes Nicolaus Brønsteds definition of acids and bases
The Dane Brønsted worked in 1923 from the idea that acids could give off H+, to a base, which could absorb this, i.e.
Acids and bases were connected, in the manner that for an acid, there was a corresponding base, and vice versa, so when the acid HCl dissociated, Cl was the corresponding base, even if Cl was not alkaline in the sense of leading to an increase in pH for the solution.

Independently, and at the same time as Brønsted, the Englishman Thomas Martin Lowry, worked on the same idea. Therefore, you can occasionally see Brønsted acids and bases referred to as Brønsted-Lowry acids and bases.

Brønsted's definition of acids and baser is the most common definition, in connection with chemistry classes.


Gilbert Newton Lewis' definition of acids and bases
The American Lewis came up with another way of looking at acids and bases in 1923. Instead of looking at the absorption and release of H+, he looked at the ability to form bonds to lon pairs, so
The model was able to explain both the reactions of Brønsted acids/bases and a number of elements' and chemical compound's ability to cause acidic or alkaline conditions, despite being unable to give off or take up H+. It could be something like the reaction explaining why CO2 acidifies water, a currently very relevant reaction in connection with the problem with acidification of the seas:

CO2(g) + H2O(l) H2CO3(aq)
H2CO3(aq) 2 H+(aq) + CO32−(aq)

or why Al3+ ions leads to an acidic solution:

Al3+(aq) + 6 H2O(l) Al[H2O]63+
Al[H2O]63+(aq) + H2O(l) Al[H2O]5[OH]2+(aq) + H3O+(aq)


Acids are divided into strong acids, medium strong acids, weak acids and very weak acids. The classification is based on the property called the acid constant. For bases you have the same classification, based on the base constant. The constants are independent of the definitions for acids and bases. The exact nature of these constants is shown under acid/base equilibria.

Besides the classification based in the acid constant, acids and bases are also grouped as organic and inorganic. The difference is mostly important for handling, storage and application areas. For the H+ ion in solutions it makes no difference whether it came from an organic or inorganic compound.

One of the more important classifications is the number of H+ that can be given off or absorbed per molecule. For acids you refer to them as protic or valent, and for baser you refer to them as basic. For an acid that can give off 1 H+, you refer to them as monoprotic or monovalent acid. If they can give off 2 or 3 H+, they are referred to as diprotic/divalent and triprotic/trivalent acid respectively. Similarly for bases, they are referred to as mono-, di- and tribasic when they can absorb 1, 2 or 3 H+, respectively. How this looks and works in practice, we will take a closer look at in acid/base equilibria.

One important detail: It is correct that H+ is given off and absorbed, but in aqueous solution, the following reaction takes place:

H2O(l) + H+(aq) H3O+(aq)

H3O+ is called the oxonium ion, and that is the actual way that H+ is present in the water. When writing chemical equations, it is customary to write H+(aq), implying that it is actually present as the oxonium ion, but it is only H+ which is relevant.

Can you have chemical compounds that are both acids and bases, then? Yes, you can. They are called ampholytes. In principle they can also be called zwitter ions, but zwitter ions are only ions having both positive and negative charges and do not have to be related to acids/bases. Ampholyte refers directly to being able to act as both acid and base.




pH and pOH

Most people have heard about pH, and many people also have an idea that it has something to do with how much the acid or solution can corrode. This is bot correct and incorrect. pH is a measure for the actual concentration of H+, [H+], and often you will see that more H+ will cause more corrosion and vice versa, but acids and bases are not the only corrosive materials, and you have several factors influencing whether a certain pH value will cause a corrosion or not.

pH is an abbreviation of potentia hydrogenii, i.e. the potential for hydrogen. The syntax here carries information, so it is written pH and not ph, Ph or PH, not even at the beginning of a sentence, even though English orthography normally specifies otherwise.

Solutions having pH = 7 are normally referred to as neutral. At pH < 7 you refer to them as acidic solutions, and at pH > 7 you call them basic or alkaline solutions.

pOH (potential hydroxide) corresponds to pH, only here it is the actual concentration of OH, [OH], in solution. I can seem a bit odd to measure OH, when not all bases contain OH. There is a perfectly good explanation for that, and it will be addressed under acid/base equilibria. pOH as a value being declared, is not used very often, except for exercises in chemistry classes. It is used widely in calculations on bases, so it is not like pOH has no use, quite the opposite, it is just not the value being shown. It is converted to the pH value instead.

Now, starting with the definitions, they look like this:

pH: -log[H+]
pOH: -log[OH]

Besides this, we also have that:

pH + pOH = 14

This is an effect of water's autoprotolysis, which is addressed as a part of acid/base equilibria, and the reason why we still use pH in alkaline solutions. You would think that for solutions containing OH it only made sense talking about pOH; pH is a referral to the concentration of acid, and here we have a base. This is where water's autoprotolysis comes in, and makes it possible.




The pH scale

It is a common misunderstanding that the pH scale goes from 1 to 14. One of the reasons for this, is most likely that indicator paper for pH measurements ranges from 1 to 14, so it is a reasonable assumption that the scale does not extend beyond that. This, however, is not correct.

There is no problem getting pH values less than 1, you can even have negative values. Likewise, pH can easily be more than 14. In principle, there is no limit to how high or low pH can be, but because there is a limitation to the solubility of acids and bases, there is also a natural limit for pH values. Concentrated hydrochloric acid (HCl) e.g. is 37.2 % (by weight) or 10.2 mol/l, which is equivalent to pH -1, and concentrated NaOH is 50.5 % (by weight) or 19.4 mol/l, which is pH 15,3.

pOH values can, similar to pH, also be less than 1 and more than 14, with the same possibilities and limitations as for pH.