Perchlorate [ClO4]- ion Lewis dot structure, molecular geometry or shape, electron geometry, bond angle, hybridization, formal charges, polar vs. non-polar
ClO4– is the chemical formula for the perchlorate ion. It comprises a chlorine atom which is bonded to four atoms of oxygen. The root anion for ClO4– is the chlorate (ClO3–) ion. Per is added as a prefix to the name chlorate due to one excess oxygen added to the root anion.
In this article, we will discuss some interesting and extremely valuable information about the perchlorate [ClO4]– ion, including how to draw its Lewis dot structure, what is its molecular geometry or shape, electron geometry, bond angle, hybridization, formal charges, polarity, etc.
So all of you, chemistry students, dive into the article and start reading to expand the horizons of your chemistry-related knowledge.
The Lewis structure of a perchlorate [ClO4]– ion consists of a chlorine (Cl) atom at the center; it is bonded to four atoms of oxygen (O) at the sides. There are a total of four electron density regions around the central Cl atom in [ClO4]– lewis structure.
All 4 electron density regions are constituted of bond pairs; thus, there is no lone pair of electrons on the central Cl atom in the ClO4–.
Drawing the Lewis dot structure of perchlorate [ClO4]– ion is easy if you follow the following simple steps.
Steps for drawing the Lewis dot structure of [ClO4]–
1. Count the total valence electrons in [ClO4]–
The very first step while drawing the Lewis structure of [ClO4]– is to count the total valence electrons present in the concerned elemental atoms.
There are two different elemental atoms present in the perchlorate ion i.e., a chlorine (Cl) atom and an oxygen (O) atom.
Oxygen is present in Group VI-A of the Periodic Table so it has a total of 6 valence electrons. Conversely, chlorine is present in Group VII-A so it has a total of 7 valence electrons in each atom.
The [ClO4]– ion consists of 1 Cl-atom, 4 O-atoms and it also carries a negative (-1) charge which means 1 extra valence electron. Thus, the valence electrons in the Lewis structure of [ClO4]– = 1(7) + 4(6) + 1 = 32 valence electrons.
2. Choose the central atom
In the second step of drawing the Lewis structure of a molecule or a molecular ion, we need to place the least electronegative atom at the center.
As electronegativity refers to the ability of an elemental atom to attract a shared pair of electrons from a covalent chemical bond therefore the least electronegative atom is the one that is most likely to share its electrons with other atoms in its surroundings.
As chlorine (Cl) is less electronegative than oxygen (O) so the chlorine atom is placed at the center of the [ClO4]– Lewis structure while the oxygen atoms are spread around it, as shown in the figure below.
3. Connect outer atoms with the central atom
In this step, we need to join the central Cl atom with the outer O atoms using single straight lines.
Each straight line represents a single covalent bond, i.e., a bond pair containing 2 electrons. There are a total of 4 single bonds in the above diagram, which means a total of 4(2) = 8 valence electrons are used till this step, out of the 32 initially available.
Total valence electrons available – electrons used till step 3 = 32 – 8 = 24 valence electrons.
This means we still have 24 valence electrons to be accommodated in the Lewis dot structure of [ClO4]–.
4. Complete the octet of outer atoms
As we already identified, the four oxygen atoms act as outer atoms in [ClO4]–, and each O atom needs a total of 8 valence electrons in order to achieve a stable octet electronic configuration.
Single Cl-O bonds on each side of the [ClO4]– means all the outer O atoms already have 2 electrons each. Thus, all four O atoms require 6 more electrons to complete their octet.
Thus, these 6 electrons are placed as 3 lone pairs around each outer O atom in [ClO4]– Lewis structure, as shown in the figure below.
5. Complete the octet of the central atom
Total valence electrons used till step 4 = 4 single bonds + 4 (electrons placed around each O atom, shown as dots) = 4(2) + 4(6) = 32 valence electrons.
Total valence electrons available – electrons used till step 4 = 32 – 32= 0 valence electrons.
As all the valence electrons initially available are already consumed, so there is no lone pair of electrons on the central Cl-atom in the ClO4– Lewis structure.
Also, the central Cl-atom has a total of 4 single bonds around it which denotes it has 8 valence electrons hence a complete octet electronic configuration.
But is this structure stable? Let us check that using the formal charge concept.
6. Check the stability of the ClO4– Lewis structure using the formal charge concept
The less the formal charge on the atoms of a molecule, the better the stability of its Lewis structure.
The formal charge can be calculated using the formula given below.
The above calculation shows that the central Cl atom contains a +3 formal charge while each outer O atom contains a -1 formal charge.
But, as we already discussed the fewer the formal charges present on bonded atoms, the greater the stability of a Lewis structure.
Hence we can minimize the formal charges by converting lone pairs into covalent bonds; let us explain to you how that is done in the next step.
7. Minimize the formal charges by converting lone pairs into covalent bonds until a stable Lewis’s structure is obtained
A +3 formal charge on the central Cl-atom represents a deficiency of 3 electrons. So 3 new bonds with the central Cl-atom can fulfill this deficiency.
Thus, 3 lone pairs of electrons present on any three of the four outer O-atoms are now converted into covalent bonds between the central Cl atom and the concerned O-atoms, as shown below.
In this way, the central Cl atom has a total of 1 single bond and 3 double covalent bonds. In other words, there is a single-bonded O-atom and three double-bonded O-atoms in this Lewis structure.
Now we can again check its stability using the formal charge concept.
For chlorine atom
Valence electrons of chlorine = 7
Bonding electrons = 1 single bond + 3 double bonds = 2 + 3(4) = 14 electrons
Non-bonding electrons = no lone pair = 0 electrons
Thus, the formal charges present on the central Cl atom and three outer O-atoms are reduced to zero. However, a -1 formal charge is present on the fourth O-atom, which accounts for the charge present on the perchlorate ion.
Finally, the perchlorate (ClO4–) Lewis structure is enclosed in square brackets, and a -1 formal charge is placed at the top right corner, as shown below.
The above Lewis structure shows that there are a total of 14 valence electrons around the central Cl-atom which denotes an expanded octet. However, it is quite possible because the chlorine atom has a 3d atomic orbital so it can accommodate more than 8 valence electrons during chemical bonding.
Another important point is that the actual structure of the perchlorate ion is a hybrid of the resonance structures given below. Each resonance structure is a way of representing the Lewis structure of a molecule or a molecular ion.
The resonance structures show that double bonds can be formed between the central Cl-atom and any three outer O-atoms in [ClO4]–.
What are the electron and molecular geometry of ClO4-?
The molecular geometry or shape of the perchlorate [ClO4]– ion is identical to its ideal electron pair geometry, i.e., tetrahedral. As there is no lone pair of electrons present on the central Cl atom in ClO4– thus there is no distortion present in its shape and/or geometry.
Molecular geometry of [ClO4]–
The molecular geometry or shape of the perchlorate [ClO4]– ion is tetrahedral.
The absence of any lone pair of electrons on the central Cl-atom in ClO4– means there are no lone pair-lone pair and lone pair-bond pair electronic repulsions present in it.
A bond pair-bond pair repulsive effect exists, which makes the bonded electron pairs occupy the four corners of a tetrahedron, as shown in the figure below.
Electron geometry of [ClO4]–
According to the valence shell electron pair repulsion (VSEPR) theory of chemical bonding, the ideal electron geometry of a molecule containing a total of 4 electron density regions around the central atom is tetrahedral.
In ClO4–, there are 4 single bonds around the central chlorine atom which makes a total of 4 electron density regions. Thus, its electron geometry is also tetrahedral.
A more straightforward way of determining the shape and geometry of a molecule or molecular ion is to use the AXN method.
AXN is a simple formula to represent the number of atoms bonded to the central atom in a molecule or molecular ion and the number of lone pairs present on it.
It is used to predict the shape and geometry of a molecule or molecular ion based on the VSEPR concept.
AXN notation for ClO4–
A in the AXN formula represents the central atom. In ClO4–, chlorine (Cl) acts as the central atom, so A = Cl.
X denotes the atoms bonded to the central atom. 4 oxygen (O) atoms are bonded to the central Cl atom in ClO4– thus X=4.
N stands for the lone pairs present on the central atom. As no lone pairs of electrons are present on the central chlorine atom in ClO4– thus N=0.
Hence, the AXN generic formula for ClO4– is AX4N0 or AX4.
Now have a quick look at the VSEPR chart given below to identify where you find AX4.
The VSEPR chart given above confirms that molecules or molecular ions with an AX4 generic formula have an identical electron geometry and molecular geometry or shape, i.e., tetrahedral, as we already noted down for ClO4–.
Hybridization of [ClO4]–
The central chlorine (Cl) atom in perchlorate [ClO4]– is sp3 hybridized.
The electronic configuration of chlorine is 1s2 2s2 2p6 3s2 3p5.
During chemical bonding, one 3s electron and two 3p electrons of chlorine shift to three empty 3d atomic orbitals. This leaves behind a half-filled 3s orbital and three half-filled 3p orbitals. The 3s orbital of chlorine hybridizes with the 3p atomic orbitals to yield four sp3 hybrid orbitals.
Each sp3 hybrid orbital is equivalent and contains a single electron only. It possesses a 25% s character and a 75% p-character.
These sp3 hybrid orbitals of chlorine consequently form four Cl-O sigma (σ) bonds by overlapping with the sp2 and p orbitals of the four oxygen atoms.
The unhybridized d atomic orbitals of chlorine form the required pi (π) bonds by overlapping with the p-orbitals of the oxygen atoms in Cl=O double bonds.
A short trick for finding the hybridization present in a molecule or molecular ion is to memorize the table given below. You can find the steric number of a molecular ion and use that against this table to find its hybridization.
The steric number of central Cl in ClO4– is 4, so it has sp3 hybridization.
Steric number
Hybridization
2
sp
3
sp2
4
sp3
5
sp3d
6
sp3d2
The [ClO4]– bond angle
The bonded O=Cl-O atoms form an ideal bond angle of 109.5° in the symmetrical tetrahedral shape of the perchlorate [ClO4]– ion. The Cl-O bond length is 144 pm.
Although, a Cl=O double bond is expected to be stronger and shorter in length as compared to the Cl-O single bond. But, the resonance present in the molecular ion leads to four equal Cl-O bond lengths in ClO4–.
Each Cl-O bond present in the [ClO4]– ion is slightly polar due to an electronegativity difference of 0.28 units between the covalently bonded chlorine (E.N = 3.16) and oxygen (E.N = 3.44) atoms. The Cl-O electron cloud stays largely towards the O atom instead of the Cl atom.
Consequently, each O atom gains a partial negative (δ–) charge while the central Cl atom attains a partial positive (δ+) charge in the perchlorate ion. Each Cl-O bond thus possesses a specific dipole moment value (symbol µ).
However, it is due to the symmetrical tetrahedral shape of the molecular ion that the dipole moments of an upwards-pointing Cl-O bond cancels with the dipole moment of three downwards-pointing Cl-O and Cl=O bonds, respectively.
The electron cloud stays uniformly distributed overall. Thus, perchlorate [ClO4]– is overall non-polar (µ = 0).
The Lewis structure for the perchlorate [ClO4]– ion displays a total of 32 valence electrons, i.e., 32/2 = 16 electron pairs.
Out of the 16 electron pairs, there are 7 bond pairs and 9 lone pairs of electrons.
A chlorine (Cl) atom is present at the center.
It is double-bonded to three O-atoms and single-bonded to one O-atom.
Out of the 9 lone pairs, 2 lone pairs of electrons are present on each double-bonded O-atom, while 3 lone pairs are present on the single-bonded O-atom.
There is no lone pair of electrons on the central Cl atom in the ClO4– Lewis structure.
How many lone pairs are found in the Lewis structure for the [ClO4–]?
There are a total of 9 lone pairs in the ClO4– Lewis structure. 3 lone pairs of electrons are present on the single-bonded O-atom, while 2 lone pairs are present on each double-bonded O-atom in ClO4–.
What is the shape of the perchlorate (ClO4–) ion based on the VSEPR theory?
The perchlorate (ClO4–) ion has an AX4 generic formula according to the VSEPR theory. A chlorine (Cl) atom is present at the center. It is bonded to 4 atoms of oxygen (O) at the sides.
There is no lone pair of electrons on the central Cl-atom; thus, ClO4– has a tetrahedral molecular geometry or shapes identical to its ideal electron pair geometry.
What is the charge of ClO4–?
Zero formal charges are present on the central Cl-atom and three of the four bonded O-atoms.
A -1 formal charge is present on the single bonded O-atom, which is also the charge present on the perchlorate [ClO4]– ion overall.
The total number of valence electrons available for drawing perchlorate [ClO4]– ion Lewis structure is 32.
The negative 1 charge present on the perchlorate ion accounts for 1 extra valence electron in its Lewis structure.
The [ClO4]– ion has an identical electron and molecular geometry or shape, i.e., tetrahedral.
The O=Cl-O bond angle is 109.5°, while all the Cl-O bond lengths are 144 pm in [ClO4]–.
It is due to the resonance present in the perchlorate [ClO4]– ion that each Cl-O bond length is equivalent as opposed to three shorter Cl=O bonds and one longer Cl-O bond, as expected.
The central Cl-atom in [ClO4]– ion is sp3 hybridized.
The [ClO4]– ion is overall non-polar. A definite plane of symmetry is present in the molecular ion, so the dipole moments of the Cl-O bonds cancel out one another’s effect. Thus, the electron cloud stays uniformly distributed l (net µ = 0).
-1 formal charge is present on a single bonded O-atom in the [ClO4]– Lewis structure which makes the molecular ion occupy an overall charge of -1.
In conclusion, [ClO4]– is a monovalentpolyatomic anion.
Vishal Goyal is the founder of Topblogtenz, a comprehensive resource for students seeking guidance and support in their chemistry studies. He holds a degree in B.Tech (Chemical Engineering) and has four years of experience as a chemistry tutor. The team at Topblogtenz includes experts like experienced researchers, professors, and educators, with the goal of making complex subjects like chemistry accessible and understandable for all. A passion for sharing knowledge and a love for chemistry and science drives the team behind the website. Let's connect through LinkedIn: https://www.linkedin.com/in/vishal-goyal-2926a122b/
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