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Acid Base Reaction

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Acids and bases are two most common compounds which play important role in various chemical reactions. It is not much difficult to define acids and bases. There are several theories which are used to define acids and bases in different manners. For example Arrhenius definition used to concept of hydrogen ions and hydroxide ions in aqueous solutions. 


On the contrary, the Bronsted-Lowry concept is based on the transfer of protons or hydrogen ion. This theory also provided the concept of conjugated acid base pair. The latest concept is known as Lewis theory which states that an acid can accept the pair of electron while the base can donate the pair of electron. Both acids and bases are also different in their chemical and physical properties. Let’s have a look at on some common properties of acids and bases. Acids are corrosive in nature and sour in taste. They can turn the phenolphthalein solution colourless and change blue to red litmus paper.

Definition

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"The reaction of acid and base leads to the formation of salt and water. This kind of displacement reaction is called as acid-base reaction." It states that it is a reaction between acid and base to form salt and water. It is also called as neutralization reaction. We can explain these reactions with any of the acid-base theory. 

The pH of these compounds is always less than 7 and they can produce hydrogen gas when react with metals. They also release carbon dioxide with carbonates. Some common examples of acids are citric acid, lactic acid, Malic acid, hydrochloric acid and acetic acid.  Bases have slippery feel and a bitter taste. The change the colour of phenolphthalein pink and turns the red litmus paper blue. The aqueous solution of bases is called as alkalis. Alkalis are found in soap, toothpaste, cleaning agents and limewater.  

Arrhenius Definition

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The Arrhenius definition of acid and base was purposed by Swedish chemist Svante Arrhenius in 1884. His theory was based on the concept of ionization of acids and bases in their aqueous solutions.  An Arrhenius acid releases hydrogen ions in its aqueous solution and these hydrogen ions further form hydronium ions $(H_{3}O^{+})$ with water. For example the ionization of hydrochloric acid in water can be shown as below.

$HCl_{aq)} \to H_3O^+_{(aq)}+Cl^{-}_{(aq)}$
                   
In case of strong acids, the concentration of hydrogen ions will be more while in case of weak acids, it will be less because weak acids cannot dissociate completely. Due to incomplete dissociation, the un-dissociate molecules remain in equilibrium with hydrogen ions and anions. Some common examples of strong acids are hydrochloric acid, nitric acid, sulphuric acid, hydrobromic acid, hydrochloric acid and Perchloric acid whereas formic acid, acetic acids and other organic acids are good examples of weak acids. Unlike Arrhenius acids, the Arrhenius bases add hydroxide ions in their aqueous solutions. For example, sodium hydroxide forms sodium ion and hydroxide ion in its aqueous solution. 

Lewis Acid Base Reaction

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The Lewis theory was based on the transfer of electrons. According to this theory, acids act as electron pair acceptors whereas bases are an electron pair donor. In other words, there must be an electron deficiency or positive charge on Lewis acid whereas a Lewis base must contain an extra pair of electrons or lone pairs. For example there is one lone pair of electrons in ammonia therefore it acts as a Lewis base whereas BF3 is an electron deficient molecule as the octet of boron is incomplete in this molecule. The acid base reaction of ammonia and boron trifluoride forms an adduct as given below.

$NH_{3(g)}+BF_{3(g)} \to H_3N-BF_{3(g)}$

In the Lewis acid-base reaction, there is no transfer of hydrogen or hydroxide ion, but only electrons move from one to another molecule. Here the lone pair of ammonia transfers to boron trifluoride molecule.  Hence ammonia acts as Lewis base and boron trifluoride acts as Lewis acid in this neutralization reaction. Overall a neutralization reaction is the combination reaction of hydrogen ion and hydroxide ion to form water, if there are hydrogen ions and hydroxide ions in the acid and base. 

Bronsted Acid Base Reaction

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The Bronsted-Lowry concept was helpful in this concept which was purposed by British chemists Johannes Nicolaus Bronsted and Thomas Martin Lowry in 1923. According to this concept, acids are proton donors whereas bases are proton acceptors. This concept overcame the limitation of aqueous solutions and the presence of hydroxide ions. There are many compounds which can act as acid and base. Such compounds are called as amphoteric compounds. 

For example water acts as an amphoteric substance in acid base reactions. This concept was helpful for the explanation of the basic nature of ammonia in neutralisation reactions. The acid base reaction of ammonia with hydrochloric acid results the formation of ammonium chloride. In this reaction, ammonia acts as a base as it accepts hydrogen ions from HCl to NH4Cl. This concept also purposed conjugated acid-base pair. The conjugated pair of an acid or base shows just opposite behaviour. In other words, the conjugated pair of a strong acid would be a weak base. For example HCl is a strong acid therefore its conjugated base; chloride ion (Cl-) is a weak conjugated base.Similarly the conjugated acid of ammonia (base) is ammonium ion (NH4+

Name and Formula of Acids   Conjugated Base 
Perchloric Acid HClO4 ClO42-
Hydrochloric Acid HCl Cl-
Sulphuric Acid H2SO4 SO42-
Nitric Acid HNO3 NO3- 
Sulphurous Acid H2SO3 SO32-
 Phosphoric Acid H3PO4 PO43- 
 Hydrofluoric acid HF F-
Formic Acid HCOOH HCOO-
Acetic Acid CH3COOH CH3COO-
Hypochlorous Acid HClO ClO-
Hydrocyanic Acid HCN CN-

Solvent System Definition

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The solvent system concept has been used extensively as a method of classifying solvolysis reactions. For example, one can compare the hydrolysis of nonmetal halides with their solvolysis by non aqueous solvents.

3H2O + OPCl3 $\rightarrow$ OP(OH)3 + 3HCl$\uparrow$

When the ionic species formed in solution are known, the solvent system approach can be useful. n solvents that are not conducive to ion formation and for which little or nothing is about known of the nature or even the existence of ions, one must be cautions. Our familiarity with aqueous solutions of high permittivity characterized by ionic reactions tends to prejudice us toward parallels in other solvents and thus tempts us to overextend the solvent system concept.

Conjugate Acid-Base Pairs

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The base that remains when an acid donates a proton is known as he conjugate base of the acid. Likewise the acid that is formed when the base accepts a proton is known as the conjugate acid of the base.

Let’s discuss a few examples of acid base reactions. The reaction of hydrochloric acid with sodium hydroxide solution forms sodium chloride with water. Similarly the reaction of HCl with ammonia solution form ammonium chloride. In both of these reactions, we white salts. The chemical equations can be written as given below.

$NaOH_{(aq)}+HCl_{(aq)} \to NaCl_{(aq)}+H_2O_{(l)}$

$NH_{3(aq)} + HCl_{(aq)} \to NH_{4} Cl_{(aq)}$

First reaction can be explained with the help of Arrhenius concept that hydrogen ions of HCl react with hydroxide ions of NaOH and form water with sodium chloride. The reaction of ammonia with hydrochloric acid is a strong acid weak base reaction because ammonia is weakly alkaline in nature whereas HCl is a strong acid. But the reaction of HCl with ammonia cannot be explained with the help of this concept as there is no hydroxide ion in ammonia. Another example of such acid-base reactions is the reaction of sulphuric acid with ammonium hydroxide.

$H_2SO_4 + 2NH_4OH \to (NH_4)_2SO_4 + 2H_2O$

Acid Alkali Reaction

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In the neutralisation reactions, there must be some way to identify the acidic or basic nature of salts. The combination of a strong acid and a strong base has always forms a neutral salt. For example, acid base reaction of potassium hydroxide and hydrochloric acid forms potassium chloride which is a neutral salt because both the constituent acids and bases are strong.

$HCl_{(aq)}$ + $KOH_{(aq)}$ $\to$ $H_2O_(l)$ + $KCl_{(aq)}$

In any salt, the cation comes from the base and anion comes from acid. The combination of a strong acid and a weak base results the formation of an acidic salt such as acid-base reaction of hydrochloric acid with ammonium hydroxide forms an acidic salt, ammonium chloride. Similarly the combination of strong base and a weak acid forms a basic salt. For example the reaction of sodium hydroxide with acetic acid forms sodium acetate which is alkaline in nature. 

Some other examples of salts with their constituent acids and bases are as given below.

Salt Base Acid Nature
$Fe(NO_3)_3$
Iron (III) hydroxide
(Weak base) 
Nitric acid (Strong acid) Acidic
$MgSO_4$ Magnesium hydroxide (Strong base)  Sulphuric acid (Strong acid) Neutral
$Ni(ClO_4)_2$
Nickel (II) hydroxide 
(Weak base)
Perchloric acid (Strong acid) Acidic
$CuSO_4$
Copper (II) hydroxide
(Weak base)
Sulphuric acid (Strong acid) Acidic
$CH_3COOH$
Acetic acid 
(Weak base)
Potassium hydroxide (Strong base) Basic 
$MgCl_2$ Magnesium hydroxide (Strong base) Hydrochloric acid (Strong acid) Neutral


Salts also play an important role in living bodies. When we suffer from acidity in the stomach we usually take medicines. Do you know this acidity problem is due to the presence of excess of gastric juice which contains hydrochloric acid? Excess of hydrochloric acid causes acidity which can be neutralized with the help of any base that is found in medicine. Usually magnesium hydroxide or milk of magnesia is used to neutralise the effect of gastric juice.  The acidic or basic nature of the substance can be determined with the help of the pH scale. The pH scale is the negative logarithm of the concentration of hydrogen ions present in the solution.  Similarly the concentration of hydroxide ions can be indicated with the help of pOH which is the negative logarithm of the concentration of hydroxide ions of a base. The addition of pH and pOH is always equal to 14. The pH of pure water is 7 while as the acidity increases, the pH of the solution decreases while basic solutions have high pH values.

Examples

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Let’s observe the acid base reaction with the help of an experiment. Take some lime solution in a beaker and add a few drops of phenolphthalein to the solution. Since the lime solution is alkaline in nature therefore the solution turns pink in colour as phenolphthalein shows pink colour in basic medium.  Now take a straw and blow some air into this pink solution. The pink colour of solution fades. Do you know why it happens? This is because as we blow air into the pink solution, carbon dioxide gas gets mixed in the solution.  Carbon dioxide gas is acidic in nature and lime solution is alkaline, therefore both involve in acid-base reaction to form salt and water. 

The reaction can be written as given below.

CO+ H2O $\to$ H2CO3

Ca(OH)2 + H2CO3 $\to$ CaCO3 + 2HOH

First carbon dioxide reacts with the water and form carbonic acid which is a weak acid. This weak acid reacts with lime solution that is calcium hydride, and forms calcium carbonate and water. Due to the presence of acid, alkaline lime solution gets neutralized and turns the phenolphthalein colorless. Salts can also prepare with the reaction of acids with metals. Here metals act as a base as they can give electrons as per Lewis theory. One of the most common examples is the reaction of hydrochloric acid (HCl) with zinc granules which results the formation of zinc chloride and liberates hydrogen gas.

$2HCl$  + $Zn$ $\to$ $ZnCl_2$  + $H_2$

Another example is the formation of magnesium sulphate from magnesium metal and sulphuric acid. 

$H_2SO_4$ + $Mg$ $\to$ $MgSO_4$ + $H2$

Or acetic acid (a weak acid) reacts with calcium to form the calcium acetate (salt) with the liberation of hydrogen gas.

 $2CH_3COOH$ + $Ca$  $\to$  $Ca(OOCCH_3)_2$ + $H_2$

Practice Problems

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Let’s do some problems of neutralisation reactions. 

Solved Examples

Question 1: Identify the reactants and products in the blanks.

  • HCl + NaOH  ______________+ $H_2O$
  • $H_2SO_4 + 2$ ______________  $(NH_4)2SO_4 + 2 H_2O$
  • 2 ______________+ $H_2CO_3  Na_2CO_3 + 2 H_2O$
  • $CO_2 + H_2O$  ___________________
  • $Ca(OH)_2 + H_2CO_3$  __________________ +$ 2 H_2O$ 
  • $Mg(OH)_2 + HCl$  ___________________ + $H_2O$
  • $H_2SO_4  + 2NH_4OH$ _______________+ $2H_2O$

Solution:
  • $HCl + NaOH \to  NaCl + H_2O$
  • $H_2SO_4 + 2 NH_4OH  \to (NH_4)2SO_4 + 2 H_2O$
  • $2 NaOH + H_2CO_3 \to Na_2CO_3 + 2 H_2O$
  • $CO_2 + H_2O \to H_2CO_3$
  • $Ca(OH)_2 + H_2CO_3 \to CaCO_3 + 2 H_2O$ 
  • $Mg(OH)_2 + HCl \to MgCl_2 + H_2O$
  • $H_2SO_4  + 2NH_4OH \to (NH_4)2SO_4 + 2H_2O$


Question 2: Identify the conjugated acid or base of the given compounds.

 Conjugate Base Acid  Base  Conjugate Acid 
   $HClO_2$  $H_2O$  
   $H_2O$  $OCl^-$  
   HCl  $H_2PO_4^-$  


Solution:
Solution:

Conjugate Base  Acid  Base  Conjugate Acid 
 $ClO_2^-$  $HClO_2$  $H_2O$  $H_3O^+$
 $OH^-$  $H_2O$  $OCl^-$  $HOCl$   
 $Cl^-$  $HCl$  $H_2PO4^-$  $H_3PO_4$