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Acids and Bases – A Brief Overview!

Acids, bases and neutralisation reactions are perhaps among the first names you hear as you enter the world of chemistry. The wonder you had as junior students when you came to know two chemicals mix to form something you regularly use in your meal for taste are among the first memories for a science student, as they started learning this amazing subject. Lemons, tamarind, soaps, baking soda and what not, we make use of acids and bases so much in our daily life. Let’s have a deeper understanding of acids and bases through this article.

What are acids and bases?

One of the primitive definitions to acids is that they are sour in taste, and change the colour of blue litmus to red. Whereas alkali – a more general term for bases, are those which are bitter, slippery to touch and change the colour of red litmus to blue. But it is not possible every time to taste and categorise chemicals and the change of colour of litmus serves as just a method to identify acidic or alkaline nature. So there is a need for a more general definition that explains the characteristics, properties, and behaviour of acids and alkalies.

An acid and alkali react with each other to form water and a salt. Another common behaviour of acids and alkalies are that they liberate H2 gas on reaction with metals. We use many theories to characterise acids and alkalies at present. The most accepted among them are:

  1. Arrhenius concept of acids and alkalies
  2. Bronsted-Lowry concept
  3. Lewis acids and bases

The Arrhenius concept

Arrhenius acids are those which increase the concentration of H+ ions in their aqueous state. That is, a substance qualifies as an acid if it ionises in water to furnish H+ ions, thereby increasing the concentration of H+ in the solution.

HCl(aq) → H+(aq) + Cl (aq)

Similarly, an alkali ionises to furnish OHions in the aqueous solution, thereby increasing the concentration OHof in the solution.

NaOH(aq) → Na + (aq) + OH (aq)

One of the common mistakes in our understanding is that a substance is acidic if it has H atoms in its formulae and basic if it has OH. However, one must refrain from such a generalisation, because there are compounds like NH3 and other amines which are basic in nature, and compounds like alcohols, for e.g C2H5OH, which are weakly acidic despite having an OH group. This might become more clear for the reader after going through the other theories mentioned here.

Bronsted-Lowry concept

This concept defines acids or bases on the basis of H+ transfer. According to this, acids are those capable of donating an H+ ion and bases are those which are capable of accepting a H+ ion.

NH 3 (g) + HCl(g) → NH 4 +Cl(s)

Here HCl donates its H+ ion to NH3. Hence NH3 acts as a base and HCl acts as an acid. Note that the mentioned reaction is a neutralisation reaction – between an acid and base to form ammonium chloride salt.

Some important facts:

  • Bronsted-Lowry concept forms a more general definition than Arrhenius concept, to include a broader spectrum of compounds under the tag acidic or basic, especially organic compounds. For e.g NH3 is identified as the base here, which wouldn’t be the case if we consider the Arrhenius concept alone.
  • Amphoteric substances: They are substances which can act as both acids and bases according to the situation. Water is a good example for this, which is illustrated in the following chemical equations :
    As base : HNO 3 (aq) + H 2 O(l) → H 3 O + (aq) + NO 3 (aq)
    As acid : NH 3 (aq) + H 2 O(l) ⇌ NH 4 + (aq) + OH (aq)
  • Acidic nature of Organic compounds: The donation or acceptance of H+ ion in organic compounds is owed to effects of other factors like lone pairs, electron cloud shift, the presence of functional groups etc. Hence, the acidic and basic behaviour is not so obvious from the presence of H or OH. The compounds or ions can be termed acidic or basic only by their behaviour as explained by the Bronsted-Lowry concept or Lewis concept.

Lewis acids and bases

Lewis concept is also very similar to the Bronsted-Lowry concept, the difference being in the fact that the transfer of lone-pair of electrons is taken as the criteria here.

Lewis acids are those which can accept a lone pair of electrons. Compounds like BF3 and AlCl3 are Lewis acids and acts as catalysts in many organic reactions.

Lewis bases are those which can donate a lone pair of electrons. Ammonia and its derivatives are common examples for Lewis bases.

pH scale

pH scale stands for ‘Hydrogen potenz’ scale and is a quantitative method to compare the relative acidic strength of species. It is the negative logarithm of H+ ion concentration.

pH = -log10[H+]

The pH scale ranges from 1 to 14. The species of pH neutral species is taken as 7. Below 7, they are acidic and above 7 are considered basic.

Acid-base systems are of great significance in chemical and biological systems. Human blood is a buffer system (systems whose pH range remains the same), thanks to the acid-base-equilibria concepts, whose understanding requires knowledge of equilibrium and ionic products which you can further read on. Natural systems like aquatic systems are very sensitive to changes in pH. Even in industry, many important chemical processes use acids and bases as catalysts and are optimised based on pH levels of the system, stressing on the significance of learning them well indeed.

There are several methods of defining acids and bases. While these definitions don’t contradict each other, they do vary in how inclusive they are. The most common definitions of acids and bases are Arrhenius acids and bases, Brønsted-Lowry acids and bases, and Lewis acids and bases. Antoine Lavoisier, Humphry Davy, and Justus Liebig also made observations regarding acids and bases, but didn’t formalize definitions.

Svante Arrhenius acids and bases

The Arrhenius theory of acids and bases dates back to 1884, building on his observation that salts, such as sodium chloride, dissociate into what he termed ions when placed into water.

  • acids produce H+ ions in aqueous solutions
  • bases produce OH ions in aqueous solutions
  • water required, so only allows for aqueous solutions
  • only protic acids are allowed; required to produce hydrogen ions
  • only hydroxide bases are allowed

Johannes Nicolaus brønsted – Thomas martin lowry acids and bases

The Brønsted or Brønsted-Lowry theory describes acid-base reactions as an acid releasing a proton and a base accepting a proton. While the acid definition is pretty much the same as that proposed by Arrhenius (a hydrogen ion is a proton), the definition of what constitutes a base is much broader.

  • acids are proton donors
  • bases are proton acceptors
  • aqueous solutions are permissible
  • bases besides hydroxides are permissible
  • only protic acids are allowed

Gilbert Newton Lewis acids and bases

The Lewis theory of acids and bases is the least restrictive model. It doesn’t deal with protons at all, but deals exclusively with electron pairs.

  • acids are electron pair acceptors
  • bases are electron pair donors
  • least restrictive of the acid-base definitions

Properties of acids and bases

Robert Boyle described the qualities of acids and bases in 1661. These characteristics may be used to easily distinguish between the two sets up chemicals without performing complicated tests:


  • taste sour (don’t taste them!)… the word ‘acid’ comes from the Latin acere, which means ‘sour’
  • acids are corrosive
  • acids change litmus (a blue vegetable dye) from blue to red
  • their aqueous (water) solutions conduct electric current (are electrolytes)
  • react with bases to form salts and water
  • evolve hydrogen gas (H2) upon reaction with an active metal (such as alkali metals, alkaline earth metals, zinc, aluminum)


  • taste bitter (don’t taste them!)
  • feel slippery or soapy (don’t arbitrarily touch them!)
  • bases don’t change the color of litmus; they can turn red (acidified) litmus back to blue
  • their aqueous (water) solutions conduct an electric current (are electrolytes)
  • react with acids to form salts and water

Examples of common acids

  • citric acid (from certain fruits and veggies, notably citrus fruits)
  • ascorbic acid (vitamin C, as from certain fruits)
  • vinegar (5% acetic acid)
  • carbonic acid (for carbonation of soft drinks)
  • lactic acid (in buttermilk)

Examples of common bases

  • detergents
  • soap
  • lye (NaOH)
  • household ammonia (aqueous)

Strong and weak acids and bases

The strength of acids and bases depends on their ability to dissociate or break into their ions in water. A strong acid or strong base completely dissociate (e.g., HCl or NaOH), while a weak acid or weak base only partially dissociates (e.g., acetic acid).

The acid dissociation constant and base dissociation constant indicates the relative strength of an acid or base. The acid dissociation constant Ka is the equilibrium constant of an acid-base dissociation:

HA + H2O ⇆ A + H3O+

where HA is the acid and A is the conjugate base.

Ka = [A][H3O+] / [HA][H2O]

This is used to calculate pKa, the logarithmic constant:

pka = – log10 Ka

The larger the pKa value, the smaller the dissociation of the acid and the weaker the acid. Strong acids have a pKa of less than -2.

Some quick facts about Acids and Bases

  • Any aqueous (water-based) liquid can be classified as an acid, base, or neutral. Oils and other non-aqueous liquids are not acids or bases.
  • There are different definitions of acids and bases, but acids can accept an electron pair or donate a hydrogen ion or a proton in a chemical reaction, while bases can donate an electron pair or accept hydrogen or a proton.
  • Acids and bases are characterized as strong or weak. A strong acid or strong base completely dissociates into its ions in water. If the compound does not completely dissociate, it’s a weak acid or base. How corrosive an acid or a base is does not relate to its strength.
  • The pH scale is a measure of the acidity or alkalinity (basicity) or a solution. The scale runs from 0 to 14, with acids having a pH less than 7, 7 being neutral, and bases having a pH higher than 7.
  • Acids and bases react with each other in what is called a neutralization reaction. The reaction produces salt and water and leaves the solution closer to a neutral pH than before.
  • One common test of whether an unknown is an acid or a base is to wet litmus paper with it. Litmus paper is a paper treated with an extract from a certain lichen that changes color according to pH. Acids turn litmus paper red, while bases turn litmus paper blue. A neutral chemical won’t change the paper’s color.
  • Because they separate into ions in water, both acids and bases conduct electricity.
  • Acids and bases are important in the human body. For example, the stomach secretes hydrochloric acid, HCl, to digest food. The pancreas secretes a fluid rich in the base bicarbonate to neutralize stomach acid before it reaches the small intestine.
  • Acids and bases react with metals. Acids release hydrogen gas when reacted with metals. Sometimes hydrogen gas is released when a base reacts with a metal, such as reacting sodium hydroxide (NaOH) and zinc. Another typical reaction between a base and a metal is a double displacement reaction, which may produce a precipitate metal hydroxide.

Chart Comparing Acids and Bases

Characteristic Acids Bases
reactivity accept electron pairs or donate hydrogen ions or protons donate electron pairs or donate hydroxide ions or electrons
pH less than 7 greater than 7
taste (don’t test unknowns this way) sour soapy or bitter
corrosivity may be corrosive may be corrosive
touch (don’t test unknowns) astringent slippery
litmus test red blue
conductivity in solution conduct electricity conduct electricity
common examples vinegar, lemon juice, sulfuric acid, hydrochloric acid, nitric acid bleach, soap, ammonia, sodium hydroxide, detergent

That’s all on acids and bases for now. You can read about s-block elements here.

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