Simple Method for writing Lewis Structures of perchloric acid HClO₄

A simple procedure for writing Lewis structures is given in a previous article entitled “Lewis Structures and the Octet Rule”. Several worked examples relevant to this procedure were given in previous posts please see the Sitemap - Table of Contents (Lewis Electron Dot Structures).

Another example  for writing Lewis structures following the above procedure is given below.

Let us consider the case of the Lewis electron dot structures of perchloric acid HClO₄. Perchloric acid is a colorless liquid. It is a stronger acid than nitric and sulfuric acid. Perchloric acid is useful for preparing ammonium perchlorate, an important rocket fuel component. It is also used as:

a solvent for metals and alloy
a dehydrating agent, particularly in the determination of silica in iron and steel and in cement and other silicate materials
an oxidizing agent, especially in the determination of chromium in steel, ferrochrome, chromite, leather
a solvent for sulfide ores for the determination of copper and other metals

How can we construct the Lewis structure of HClO₄ ?

 

Step 1: Connect the atoms with single bonds. Chlorine is the central atom:

 

Step 2: Calculate the # of electrons in π bonds (multiple bonds) using  formula (1) in the article entitled “Lewis Structures and the Octet Rule”. 

Where n in this case is 5 since HClO₄ consists of 6 atoms but one of them is a hydro gen atom (remember n is the number of atoms in a molecule minus the hydrogen atoms).

Where V = (1 + 7 + 4*6 ) = 32   

Therefore, P = 6n + 2 – V = 6 * 5 + 2 – 32 = 0

So there are no π electrons in HClO4 and therefore the structure of Step 1 is the Lewis structure.

Electrons are placed around each atom so that the octet rule is obeyed. Formal charges are assigned and equalized using resonance.

 

Step 3 & 4: The Lewis structure for HClO₄ is as follows:

Simple method for writing Lewis Structures – Ozone O3 and carbonate CO3-2

Simple Method for writing Lewis Electron Dot Structures - Ozone O₃ and carbonate CO₃^-2

A simple method for writing Lewis electron dot structures is given in a previous article entitled “Lewis Structures and the Octet Rule”. Examples for writing Lewis structures following the above procedure are given bellow:

Consider the case of ozone O₃ Lewis electron dot structures:

Step1: The central atom will be one of the oxygen atoms.  Connect the 3 atoms with a single bonds

O – O – O

             

Step 2: Calculate the # of electrons in p bonds (pi bonds, multiple bonds) using formula (1) in the article entitled “Lewis Structures and the Octet Rule”.

Where n in this case is 3 since O₃ consists of three atoms

Where V = (6 + 6 + 6 )  = 18  

Therefore, P = 6n + 2 – V = 6 * 3 + 2 – 18 = 2      Therefore, there are 2 π electrons (pi electrons) in O₃  and so 1 double   bond must be added to the structure of Step 1.

Step 3 & 4:  The 2 atoms are joined together with a double bond. Therefore the Lewis electron dot structures for O₃ are as follows:



 

Consider the case of the Lewis electron dot structures of the carbonate ion, CO₃^-2

Carbonate species act as buffers and are necessary for all biological systems. They also play important role in neutralization of strong acids and bases. The carbonate system as it is called is the major source of buffering in the ocean and in natural waters.

Step1: The central atom will be the C atom since it is the only atom with “subscript” equal to 1 in the molecular formula.  Connect the O atoms with the C atom with single bonds.

 

             

Step 2: Calculate the # of electrons in π bonds (pi bonds, multiple bonds) using  formula(1) in the article entitled “Lewis Structures and the Octet Rule”.:

Where n in this case is 4 since CO₃^-2consists of four atoms

Where V = (4 + 6 + 6 + 6 ) – (-2)  = 24   Therefore, P = 6n + 2 – V = 6 * 4 + 2 – 24 = 2       

There are 2 π electrons (pi electrons) in  CO₃^-2  and therefore 1 double  bond must be added to the structure of Step 1.

Step 3 & 4: One double bond between C and O is added to the structure in step 1. Unshared electron pairs are added so that there is an octet of electrons around each atom. All the equivalent resonance structures are drawn by delocalizing electron pairs. Therefore, the Lewis electron dot structures for CO₃^-2 are as follows:

Figure 4: Lewis structures for the carbonate ion CO₃^-2. The 3 structures are equivalent since they have equal formal charges on the elements. It has been proven experimentally that the C-O bond in the carbonate ion is a hybrid of a single and a double bond (the length of the C-O bond in the carbonate ion is approximately half of the sum of the lengths of a normal C-O and C=O bond).This means that the theoretical Lewis structures are in agreement with experimental results.