Whenever we discuss acids and bases, the first example that comes to our mind is that acidic foods taste sour such as lemon juice containing citric acid. Basic foods taste bitter such as the taste of garlic and spinach. But there is much more to the concept of acids and bases other than their sour and bitter tastes. In this article, we will introduce you to all the concepts related to acid-base chemistry. So, continue reading!
What are acids and bases
The earliest definition of an acid is that it is a chemical substance that dissolves in water to produce hydrogen (H+) ions. In contrast to that, a base is a chemical substance that releases hydroxide (OH-) ions in water or an aqueous solution.
Acids turn a blue litmus paper red, while a base does the opposite by turning a moist red litmus paper blue. Acidic substances taste sour and have a slippery texture, while basic or alkaline substances have a bitter taste and a soapy texture. On the pH scale of 0 to 14, acids have a pH value below 7, while the pH of a base lies above the neutral pH of 7. For instance, the pH of a strong acid such as hydrochloric acid (HCl) is 1.1, while a strong base such as sodium hydroxide (NaOH) has a pH level close to 14.
However, the definition given above is according to the Arrhenius concept of acids and bases, first introduced as far back as the late 1880s. But as this concept evolved over the passage of time, the acid-base definition also transformed as per new theories such as the Bronsted-Lowry theory and the Lewis theory.
Let us recall all three important theories related to the acid-base concept one by one.
Arrhenius concept of acids and bases
The history of acids and bases goes back to the first theory introduced by Svante August Arrhenius.
This theory known as the Arrhenius theory or concept of acids and bases states that acid is a substance that produces hydrogen (H+) ions in the water while a base gives hydroxide (OH–) ions in water.
However, the limitation of the Arrhenius theory is that it does not take into account the chemical substances that do not produce OH– ions but are still base in nature. For e.g., ammonia (NH3) is a base, but it does not possess OH– ions.
Bronsted-Lowry theory of acids and bases
The Bronsted-Lowry theory, introduced in 1923 by Johannes Nicolaus Bronsted (Danish Chemist) and Thomas Martin Lowry (English Chemist), defines acids and bases as proton (H+) donors and acceptors, respectively.
Bronsted –Lowry acids dissociate to release H+ ions or protons in a solution.
Bronsted-Lowry bases accept H+ ions or protons from a solution.
So, in the presence of a Bronsted-Lowry acid, the concentration of H+ ions increase in the solution. The greater the strength of an acid, the more H+ ions will be released. As pH is defined as the power of hydrogen and calculated as negative log[H+] so increase in H+ decreases the pH value, as we already noted down for strong acids.
Contrarily, the presence of a strong Bronsted-Lowry base decreases the H+ concentration in a solution, so the pH value increases. However, the pOH value of the solution decreases which is calculated as pOH + pH = 14 or pOH = 14 –pH.
pOH is also calculated as negative log [OH–], so the pOH of a strongly basic solution decreases both via the Arrhenius acid-base concept and the Bronsted-Lowry theory of acids and bases.
Conjugate acids and bases
According to the Bronsted-Lowry acid base concept, acids are proton donors, while bases being the opposite, accept protons. In this way, an acid and a base act as a conjugate acid-base pair.
For e.g., in an acid-base chemical reaction between water (H2O) and ammonia (NH3), H2O acts as a Bronsted-Lowry acid and donates a proton, while NH3 is the Bronsted-Lowry base that accepts this proton, as shown below.
In the above example, the chemical constituent formed after proton donation i.e., OH– is known as the conjugate base of acid. Conversely, the chemical entity formed as a result of accepting a proton, i.e., NH4+ is known as the conjugate acid of a base.
A strong acid gives a weak conjugate base while a weak acid gives a strong conjugate base, and likewise for the bases and their conjugate acids.
Example:
The conjugate base of sulfuric acid (H2SO4), a strong acid, is HSO4– which is a weak conjugate base.
Contrarily the conjugate base of acetic acid (CH3COOH), a weak acid, is CH3COO– which is a strong conjugate base.
The strength of an acid is also determined by its acid dissociation constant Ka. The higher the Ka value, the greater the acid strength because, in that case, it easily and readily dissociates to release H+ ions in a solution.
The acid dissociation constant Ka of a weak acid such as CH3COOH is related to its pH under equilibrium conditions by the equation given below:
pH = pKa + log {[A–]/[HA]}
where [A–] = concentration of conjugate base and [HA] = concentration of acid
pKa = – log Ka
Always remember for a strong mineral acid such as HCl, H2SO4 etc., as the acid completely dissociates to give H+ ions in the solution, so [HA] = [A–], log 1 = 0, so pH = pKa. Strong acids have a high Ka, which means an extremely small pKa value and, consequently a low pH.
In one example above, we mentioned that water (H2O) can act as a Bronsted-Lowry acid by donating a proton but what if it also accepts protons and can behave both as an acid and a base. How is that possible?
Let’s find that out in the next section.
Amphiprotic vs Amphoteric
A chemical substance that can act both as a Bronsted-Lowry acid and a Bronsted-Lowry base is known as amphiprotic. It can donate as well as accept protons in a chemical reaction. Thus, water (H2O) is an amphiprotic substance as it can both donate and accept protons, as shown in the following equations.
In contrast to amphiprotic substances, amphoteric is another term used to represent chemical substances that can act both as acids as well as bases. However, it is not necessary that they exhibit dual acid-base behavior by donating and accepting protons only. Let’s see an example:
Example: Zinc oxide (ZnO) is an amphoteric substance.
ZnO does not possess any proton nor accept protons, so it is clearly not a Bronsted-Lowry acid nor a Bronsted-Lowry base. Thus, it is not amphiprotic. But as it can still act both as an acid as well as a base, so ZnO is amphoteric in nature.
ZnO can also act as an acid by reacting with a base such as NaOH to yield salt and water.
All amphiprotic substances are amphoteric. However, not all amphoteric substances are amphiprotic in nature.
In the above example, ZnO does not contain H+ ions, but it reacts as a base possessing lone pairs of electrons. This concept gives birth to the Lewis theory of acids and bases.
The Lewis theory overcomes the limitations of both Arrhenius and Bronsted-Lowry theories and gives a different dimension to acid-base chemistry.
Lewis theory of acids and bases
The Lewis theory was introduced by Gilbert N. Lewis. According to this theory, acids are defined as chemical substances that accept lone pairs of electrons during chemical bonding while bases donate lone pairs of electrons.
As per the Lewis theory of acids and bases, acids are electrophilic (electron-deficient) species, while bases are nucleophilic (electron-rich) in nature.
Example:
Boron trifluoride (BF3) is a Lewis acid. It reacts in an acid-base reaction by accepting a lone pair of electrons from a Lewis base.
Ammonia (NH3) can act both as a Lewis base as well as a Bronsted-Lowry base. It accepts protons to form ammonium (NH4+) ions. Similarly, as it contains a lone pair of electrons on the nitrogen (N) atom so it can also act as a Lewis base by donating electrons.
See the acid-base reaction of BF3 with NH3, as shown below.
Difference between acids and bases
| Acids | Bases | |
| Arrhenius concept | Release H+ ions in water | Release OH– ions in water |
| Bronsted-Lowry theory | Proton donors | Proton acceptors |
| Lewis theory | Electron pair acceptor | Electron pair donor |
Difference between base and alkali
Alkali is defined as a water-soluble base. It is a term used for a base that essentially produce OH– ions in an aqueous solution. So, all Arrhenius bases are alkalis. In fact, all alkalis are bases but not all bases are alkalis.
How are acids and bases different from salts
Salt is the main product of an acid-base neutralization reaction. It is simply defined as an ionic compound made up of a cation other than the H+ ion and an anion other than OH– ion. When an acid reacts with a base, salt and water are formed.
Acid-base titration
In the chemistry laboratory, the unknown concentration of a base can be determined via an acid-base neutralization reaction by using an acid of known concentration. This acid-base reaction can be performed in a conical titration flask in the presence of an indicator, and the experiment is then called an acid-base titration. It is also sometimes known as neutralization titration.
Both an acid and/or a base can act as a titrant or a titrand as per experimental conditions.
Learn more about acid-base titration in our next article.
Also, check out our article, Introduction to Titration, for everything you need to know about titrimetry.
References
1. C.Harris, Daniel. 2010. Quantitative Chemical Analysis (W.H Freeman and Company ).
2. Shikha, Munjal, and Singh Aakash. 2020. ‘The Arrhenius Acid and Base Theory.’ in Singh Ambrish (ed.), Corrosion (Intech Open: Rijeka).
