Nandita Biswas

Hello.... I am Nandita Biswas. I have completed my master's in Chemistry with a specialization in organic and physical chemistry. Also, I have done two projects in chemistry- One dealing with colorimetric estimation and determination of ions in solutions. Others in Solvatochromism study fluorophores and their uses in the field of chemistry alongside their stacking properties on emission. I am working as a Research Associate Trainee in Medicinal Department. Let's connect through LinkedIn-https://www.linkedin.com/in/nandita-biswas-244b4b179

ClF3 Lewis Structure,Characteristics:13 Facts You Should Know

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clfa

Here, we shall learn how to draw ClF3 lewis dot structure, to count valence electrons, octet rule, its solubility and other such important characteristics.

ClF3 lewis structure is an inter-halogen compound that plays a very important role as solvent, in nuclear chemistry, therefore knowing ClF3 lewis structure, its bonding and connectivity with atoms is very crucial.

ClF3 lewis structure is a simple electronic representation of the skeletal structure of the molecule, about how the electrons are arranged around the atoms.

How to draw ClF3 lewis structure ?

  1. ClF3 lewis structure can be drawn by first by counting the total valence electrons of all the atoms combined. Cl has electronic configuration : [Ne]3s23p5 and F electronic configuration : [He]2s22p5. Therefore, it has a total of 28 valence electrons available.
  2. The central atom is chosen based on their electronegativity and a skeletal structure is drawn. Electronegativity of Cl = 3.16 and F =3.98, thereby choosing Cl has the central atom.
  3. Each atom tries to fulfil its octet by accommodating 8 electrons around it to follow octet rule. A single bond is drawn from each atom with the valence electrons to the nearby atoms.
  4. The Cl is surrounded by a total of 10 electrons in the Lewis dot structure, thereby, violating octet rule.It can accommodate extra electrons apart from the 8 electrons already assigned through bond and lone pairs is, because it has expanded empty 3d shells.
  5. The remaining electrons not forming covalent bond will stay as lone pair of electrons.

Note: Elements having expanded valence shells like 3d elements, it can exceed the octet rule like SF6 , PF5  or elements with fewer valence electrons can have incomplete octet like H2 .

Construction of ClF3 Lewis Structure step by step :

clfa
ClF3 Lewis structure

ClF3 lewis structure formal charge :

ClF3 lewis structure formal charge briefs about the electronic charge of each atom in a molecule based on the Lewis dot structure.

Generally, formal charge can be calculated mathematically by the formula :

Formal charge = (Number of valence electrons in a free atom of the element) – (Number of unshared electrons on the atom) – (Number of bonds to the atom)

In addition, Charge on the molecule= sum of all the formal charges.

Formal Charge of Fa , Fb , Fc  = 7-6-1 = 0 (All the F atoms are equivalent)

Formal Charge of Cl = 7- 4- 3 = 0

clf3 lewis structure
ClF3 lewis structure formal charge :

ClF3 lewis structure resonance :

In ClF3 lewis structure, all the F atoms are equivalent and they cannot form double bonds as their octet gets completed with a stable noble configuration when single bond formation with the central atom takes place.

clf3c
ClF3 lewis structure has all bonds equivalent.

ClF3 lewis structure octet rule :

Octet rule states that an atom tries to bond in a manner that allows them to take 8 electrons in their valence shell to fulfil their octet. However, many molecules with atoms that has expanded subshells can take up more than 8 electrons, thereby, violating octet rule.

Here, Cl has expanded subshell 3d orbital that is empty. It takes up 2 lone pairs of electrons and 3 bond pairs giving a total of 10 electrons in their outermost shell. Thus, it violates octet rule.

The F atoms take up 8 electrons as per octet rule as they do not have expanded subshells.

ClF3 valence electrons :

Cl electronic configuration : [Ne]3s23p5

F electronic configuration : [He]2s22p5.

Each F atom has 7 outermost electrons, there are 3 F atoms making a total of 21 valence elctcrons. Cl atom has 7 valence electrons.

Therefore, it has a total of 28 valence electons available.

ClF3 lewis structure lone pairs :

From the Lewis dot structure, it is evident that Cl has 2 lone pairs of electrons. Each F atom has 3 lone pairs of electrons.

The total lone pairs of electrons are 11 .

ClF3 hybridisation :

A simple method to calculate number of orbitals taking part in bond formation using VSEPR model is thorugh a mathematical formula given below :

Hybridisation of a molecule = ( Valence electrons of the central atom + Number of monovalent atoms attached to the central atom + Negative charge on the molecule – Positive charge on the molecule )/2

ClF3 Hybridisation = ( 7 + 3 – 0 – 0 )/2 = 5 = sp3d

The central atom is Chlorine. Its electronic configuration in ground state : [Ne]3s23p5  and the excited electronic configuration : [Ne]3s23p43d1 . There are two lone pairs of electrons, one in 3s and the other in one of the 3p subshells occupying the two of the 5 hybrid orbitals. The remaining 3 unpaired electrons, two in 3p subshells and one in 3d subshells will form bonds with the 3 F atoms.

It can be either Trigonal bipyramidal or Square pyramidal geometry.

It adopts a trigonal bipyramidal geometry (Fig: a) as the lone pairs are at 1200 to each other which prevents from steric repulsion of the lone pairs as well as the bond pairs.

clf3 lewis structure
a. Trigonal bipyramidal b. Square pyramidal

ClF3  lewis structure shape :

When we talk about shape of a molecule, we do not consider their lone pairs hybridization. Therefore, the shape is about how the bond pairs are oriented in spatial dimension which gives ClF3  a T-shaped structure.

The F atoms in axial position are not exactly at 900 but somewhat at 880 due to lone pair- bond pair repulsion.

clfg
It adopts a T-shaped planar shape

ClF3 lewis structure angle :

The angle is not exactly 1800 but less than it as there is lone pair- bond pair repulsion which exceeds the bond pair-bond pair repulsions. It has approximately an angle of 1750 . Also the F atoms in axial position are not exactly at 900 but somewhat at 880  due to lone pair- bond pair repulsion.

clfh
Bond angle is not exactly 1800

Is ClF3 acidic or basic ?

It is surprising to know this interhalogen compound can act as both Lewis acid and Lewis base.

In other words, this can be termed as amphoteric. As a result, they are mainly used for establishing a solvent system.

The reaction below shows how ClF3 acts as a base and an acid in different conditions :

AsF3 + ClF3    à [ClF2]+ [AsF4]–           Here, [ClF2]+ , a cationic species is formed, therefore, it acts as an acid.

ClF3  +  NOF  à  [NO]+ [ClF4]–             Here, [ClF4] , an anionic species is formed, therefore, it acts as a base.   

Is ClF3 ionic ?

No, it is not an ionic compound.

This AX3  form of interhalogen compounds do not have great electronegativity difference.

It is a covalent liquid with ionic percentage less than 40% as per Fajan’s rule. Cl has electronegativity of = 3.16 and F = 3.98 whose difference is not appreciable enough to consider it as an ionic compound.

clfwiki
ClF3 lewis structure shape with bond angle and bond distance from wikipedia

Is ClF3 polar or nonpolar ?

It is a polar interhalogen compound. This can be justified by considering the dipole vectors.

The compound is not perfectly T shaped so the F atoms in axial positions do not cancel each other. Furthermore, the dipoles’ vectors adds up to the axial vectors which gives a resultant vector of greater magnitude than the vector in the opposite direction.

Thus, some magnitude of dipole moment is left which makes the compound a polar one.

clfi
Dipole moments of the respective bond vectors are shown

ClF3 solubility :

ClF3 is a very reactive gas so its solubility in different solvents have to be checked properly before dissolving it in it. Few solvents in which it is soluble are listed below :

  • Benzene
  • Toluene
  • Acetic acid
  • Hexane

It is soluble in these solvents retaining its stability but at lower concentration. It explodes at higher concentration.

Is ClF3 tetrahedral ?

No, ClF3 cannot be tetrahedral. According to VSEPR model of hybridization, its geometry is trigonal bipyramidal as it is most stable in that form with less steric repulsion.

Is ClF3 linear ?

No, ClF3 is not a linear molecule.

The molecule is considered to be planar as the orientation of the central atom Cl and the surrounding atoms F forming bonds with Cl are arranged in a T-shaped manner but it is not a linear molecule. Its shape is strictly limited to T-shaped which is not linear in real sense.

Conclusion :

ClF3 is an interhalogen compound with trigonal bipyramidal geometry , sp3d hybridisation and planar T-shaped which is mainly used in solvent system.

Also Read:

7 Important Charged Ions Examples You Need To Know

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charged3

Charged ions or simply ions are any atoms or group of atoms in a molecule with a net electrical charge which can be either positive or negative due to the removal of electrons from or addition of electrons to the valence shell of the orbitals and 7 important charged ions examples, their classifications are discussed in details.

  1. H+ or H3O+
  2. Ions belonging to alkali metals and alkaline earth metals like Na+, K+ , Ca2+ , Mg2+ ,  Cs+
  3. Cations belonging to transition metal series
  4. Carbocations
  5. Carbanions
  6. Oxygen in superoxide, peroxide forms
  7. Lanthanum ions

1. H+ or H3O+ :

One of the most important charged ion examples is a hydronium ion or hydroxonium cations are produced when Arrhenius acid releases protons in aqueous medium or when a proton combines with H2O molecule to give H3O+ ions as H+  ions do not stay alone in aqueous medium. They are highly studied for their various functionality.

They have a very high conductance in solution due to Grotthuss mechanism. The ability of a molecule and to what extent it produces  H+  ions provides magnitude of their acidity.

In crystalline compounds the hydroxonium ions H3O+ , H5O2+ , H7O3+ , H9O4+ , H14O62+ have been characterized.

Structure of : H3O+  is a flat pyramid while that of H5O2+  is a kind of linear structure.

However, structures of other hydroxonium cations are quite complicated and O—H-O distances differ considerably.

Note:

Grotthuss  mechanism : It is a phenomenon by which a proton moves within the solution rapidly through H-bonding between the H+ and H2O. In other words, there is a rapid exchange of proton resulting in higher conduction.

charged ions examples
Structure of H3O+ under different conditions.

2. Ions belonging to alkali metals and alkaline earth metals like Na+, K+ , Ca2+ , Mg2+ ,  Cs+ :

These monopositive ions are formed by losing one electron from their outermost s orbitals in case of alkali metals. They do not form dipositive ions as the configuration obtained after losing its valence shell s electron is that of a noble gas configuration.

Alkaline earth metals form dipositive ions by losing two of its valence electrons. They are formed as they have low ionisation energies to attain stability. Also, their high solvation energy tendency make them stabilised thermodynamically.

Na+, K+ ions play important role in biological processes as Na+ / K+ pump ,Mg2+  is required for many enzymatic reactions like Na+-K+-ATPase, it blocks calcium for easy relaxation of muscles after contraction.  Ca2+ is used for circulating blood, relaxing nerve cells, regulating heartbeat activity etc , Cs+ is radioactive and highly toxic, they have been used in nuclear chemistry.

3. Cations belonging to transition metal series :

They display a wide range of oxidation states due to the availability of quite a few number of valence electrons in their (n-1)d and ns orbitals which are less shielded by the innermost electrons. These ions have played role as  catalysis, electrodes, electroplating, ores.

Cr3+ ions from enormous number and variety of complexes like Reinecke’s salt, this salt with Cr3+ ion was widely used to precipitate primary and secondary amines, for determination of promazine.

Cr2+ is one of the strongest reducing agents in aqueous solution. Ti2+ forms intense yellow-orange color hence they can be identified using  colorimetric determination, they are used for synthetic biomolecules delivery. Fe2+ and Fe3+ is used as a redox couple and is responsible for the oxidation of iron forming rust.

4. Carbocations :

One of the most important topics in organic chemistry. They are charged cations with +1 or more than +1 charge residing on carbon atom that is bonded to only three other atoms thereby leaving one of its tetravalency empty. It contains 6 electrons with an empty p orbital thus acquiring sp2 hybridisation with trigonal planar geometry.

Shape and hybridization of  +CR3 where R= H or any other atoms : Common example is a Methyl cation with +1 charge.

charged2
Simplest representation of a carbocation : Methyl carbocation.

5. Carbanions :

They are electrophilic in nature and can be formed by reaction with alkali and alkaline earth metals by heterolytic fission. This charged ion is extremely common and important in organic chemistry.           

It adopts a trigonal pyramidal structure similar to NH with the lone pair occupying the tetrahedral valency to avoid repulsion between the bond and lone pairs of electrons.

Shape and hybridization of  –CR3 where R= H or any other atoms : Common example is a methyl anion with -1 charge.

1 charged
Simplest representation of carbanion

6. Oxygen in superoxide, peroxide forms :

Superoxide,O2– , is an oxygen free radical responsible for many biological irregular functions. This radical anion is responsible for the slow tampering with living beings. These reactive oxygen species also play a role of action in antimalarial medicines mechanism.

It is paramagnetic and is weakly attracted to magnetic field. Peroxide,O2– , is diamagnetic as it contains paired electrons.

Bond order of superoxide ion is 1.5 with smaller bond length and forms more stable peroxide compounds with more electropositive elements, and that of peroxide’s bond order is 1 with a greater bond length.

They have wide applications in medicinal chemistry. They are used as Hydrogen peroxide. Peroxides, however, decomposes under sunlight and they have to be kept in dark areas.

charged4
1. Superoxide 2. Peroxide from wikipedia

7. Lanthanum ions :

They mostly form tripositive cations because of their Zeff value and closely spaced energy levels. They are stable hard lewis acids and reducing agents. They form wide range of complexes that are used as homogenous catalysts for industrial applications. E.g. La3+ , Ln3+ , Eu 3+ .

They are used in batteries, optoelectronics devices, superconductors, rocket fuel, nuclear energy study, transistors etc.

Classification of Charged ions :

  • Anionic charged species/anions  – These charged ions carry an overall negative charge magnitude. They are formed due to the addition of extra electrons in the valence shells which outdone the protons present in the nucleus of the system/species. Generally, Non-metals show the tendency to gain electrons, mostly down the group and across the right side of the period. Common examples : Cl–  , Br– , O2– , SO42- etc.
  • Cationic charged species/cations  – They carry an overall positive charge due to loss of electrons from their valence shells but the protons number remain the same in the nucleus. Alkali metals, alkaline earth metals, early transition elements show this tendency. Common examples : Na+ , Ca2+ , Mg2+ , Hg+ , Hg2+ , Zn2+ , NH4+ , H3O+ etc.

Note : It can be either a polyatomic ion or a monoatomic ion. The magnitude of the charge can greater or lesser than zero but never equal to zero as the whole definition of charged ions lies in the positive or negative magnitude of charge.

Conclusion :

7 most important charged ion examples have been discussed which are broadly classified as H+ or H3O+ ions, alkali metals and alkaline earth metals ions like Na+, K+ , Ca2+, transition metal series ions , Carbocations, Carbanions, superoxide, peroxide, Lanthanum ions.

I3- Lewis Structure,Geometry,Hybridization: 7 Steps (Solved)

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I3– Lewis Structure

The triiodide ion (I3⁻) consists of a linear arrangement of three iodine (I) atoms, with the central I atom bonded to two terminal I atoms. It has 7 valence electrons per I atom, plus one additional electron due to the negative charge, totaling 22 electrons. The Lewis structure shows two single I-I bonds and three lone pairs on the central I atom. The end iodine atoms each have three lone pairs. I3⁻ exhibits a linear geometry with bond angles of 180°, consistent with sp³d hybridization. The presence of the extra electron on the central iodine contributes to the ion’s stability and unique chemical properties, such as its use in starch testing as a color-changing complex.

The Lewis theory is based on octet rule which states that an atom always tend to arrange 8 electrons around themselves to acquire a stable or a noble gas configuration. There are, however, some exceptions like when a molecule is electron deficient; when it has odd number of electrons; or molecules that has extra electrons in their valence shells. E.g., BH3 ,SF6 , H2, NO etc. I3- lewis structure is one of them.

I3– Lewis Structure
I3– Lewis Structure

Methods to draw a Lewis Structure :

  1. First count the total number of valence shell electrons available for each atom.
  2. Choose the least electronegative atom as the central atom and draw the remaining atoms around the central atom and start by forming a covalent bond (a bond requires two electrons) . For most atoms, there will be a maximum of eight electrons in their valence shells to fulfill octet rule.
  3. The remaining electrons not forming covalent bond will stay as lone pair of electrons.

Note: Elements having expanded valence shells like 3d elements, it can exceed the octet rule like SF6 , PF5  or elements with fewer valence electrons can have incomplete octet like H2 .

wikipedia i3

                                                                                                    I3– structure from wikipedia

How to draw a I3– Lewis structure ? :

  1. Iodine belongs to 17th group and 5th period. It has 7 valence electrons, with expanded shells to occupy any extra electrons apart from the 8 electrons. It has 3 same I atoms with same electronegativity, therefore, choose any one as the central atom. It has a total of 22 valence electrons from 3 I atoms and a negative charge.
  2. Draw a covalent bond between each atoms with the central atom to fulfill the outer orbit. In doing so, we get 3 lone pairs of electrons on the central atom and 2 lone pairs of electrons each on the surrounding atoms. As I has empty 4d shells it can expand to accommodate extra electrons thereby violating octet rule which is generally observed for heavier elements. This gives the I3– lewis structure.
i3 1
I3– Lewis structure

I3– Lewis structure Formal Charge :

Now, we have to assign their formal charge to obtain the complete stable I3– lewis dot structure.

It briefs about the electronic charge of each atom in a molecule based on the Lewis dot structure.

Generally, formal charge can be calculated mathematically by the formula :

Formal charge = (Number of valence electrons in a free atom of the element) –  (Number of unshared electrons on the atom) – (Number of bonds to the atom)

In addition, Charge on the molecule= sum of all the formal charges.

i3 2
Lewis structure of I3–
i3 3
Each atom is assigned its respective formal charge which will give the overall charge of the ion.

Formal charge of Ia = 7-6-1 = 0 Formal charge of Ib = 7-6-2 = -1  

I3– Lewis structure resonance  :

i3 4
Resonance structure I3– Lewis structure

Dative or coordinate bonds are formed by sharing two electrons covalently by single atom to the nearest neighbouring atom.

The two I atoms at the end are similar due to resonance. It avoids forming any double bonds even though it has extra subshells to accommodate electrons is because it is most stable when it acquires a linear shape to prevent steric angular strain as the single bonds do not repel each other greatly than the single bond when present in a double bond. Therefore, this form is the most stable and likely resonance structure of I3– lewis structure.

I3– Lewis structure valence electrons :

i3- lewis structure
Bond represents bond pairs. The dots represents lone pairs.

Electronic configuration of I : [Kr]4d105s25p5. Its valence electrons are 5s25p5 which counts to a total of 7 outermost electrons. I3– Lewis structure has a total of 22 valence electrons. It has 2 bond pairs ( that formed a single covalent bond here).

Note: In actual sense of chemistry, there is a coordinate bond formation between I2 and I– ( electron pairs shared completely by iodide ion, a type of covalent bond).

I3– Lewis structure lone pairs :

I3– Lewis structure has a total of 9 lone pairs of electrons that did not participate in bond formation and residing on respective I atoms.

I3– Lewis hybridisation :

There is a simple rule or equation to be followed to find out the hybridization of molecules real quick.

Hybridisation of a molecule = ( Valence electrons of the central atom + Number of monovalent atoms attached to the central atom + Negative charge on the molecule – Positive charge on the molecule )/2

Here, I3– Lewis structure hybridization = (7+2+1)/2 = 5 i.e., sp3d.

I3– Lewis structure shape :

I3– lewis structure has a trigonal bipyramidal geometry which is well justified as the central atom contains 3 lone pairs of electrons which can stay at maximum distance with less repulsion when they occupy the equatorial position at an angle of 1200 with each other. The other two bond pairs i.e., the two I atoms occupy the apical positions. Therefore, the shape of I3– lewis structure is linear.

i36 1
The I atoms in apical position are at 1800 to each other.

I3– Lewis structure angle :

The I3– Lewis structure adopts a linear shape which has an angle of 1800 .

I3- Lewis Structure
The angle between two I atoms is 1800 .
Molecule I3– , Triiodide ion
Type Polyatomic ion
Hybridisation, Geometry sp3d , Trigonal Bipyramidal
Shape Linear
Angle 1800
Bond and Lone pairs 2 , 9
Table: Chracteristics of I3–

I3– Lewis structure octet rule :

I3– Lewis structure violates octet rule as it has expanded 4d shells which can accommodate extra electrons besides the 8 electrons required to fulfil the octet rule as evident from its lewis dot structure.

Is I3– stable ?

 Yes, triiodide ion is found to be stable.

As iodine element belongs to 5th period, it has empty 5d shells which can expand its octet to accommodate the extra electron pair provided by the iodide ion, I– to I2 . They have been found to exist in aqueous solution as well as in crystalline form with a variety of cations.

Is I3– ionic or covalent ?

Triiodide ion is a polyatomic ion with an overall negative charge in one end of the linear molecule. Hence, by default it is going to be ionic. There might be some questions regarding its polarity as its dipole each other being linear in shape, but there is always some negative charge residing due to which it has some dipole moment associated with it. Also, a polyatomic ion will never exist as a free ion in a solution , its opposite charged ions will surround it always. In solid state, there is a slight deviation from linearity giving some dipole moment.

I3 8
Even though the dipoles cancels each other, some negative charge is still left.

I3– Uses:

  1. It is widely used as a redox couple in dye synthesized solar cells widely known as DSC.
  2. It is used as a conducive materials in solar panels, batteries, electrochemical cells.
  3.  It is used as an interesting ion in electrical and magnetic materials, in host-guest compounds etc.
  4. It is used for controlled reactions, as an indicator in chemical redox reactions, for stain removal of certain dyes produced during reactions.

Conclusion :

Triiodide ion, I3- Lewis structure , is a linear polyatomic ion with sp3d hybridisation and trigonal bipyramidal geometry and acquiring a linear shape known through VSEPR model.

Also Read:

Polyatomic Ion Examples:Definition,Structure,Characteristics

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polyatomic ion

Polyatomic ions are charged species that contain more than one atom held together by covalent bonds. The polyatomic ion examples we will cover in this article are:

  • Ammonium
  • Acetate
  • Carbonate
  • Chromate
  • Cyanide
  • Hydroxide
  • Nitrate
  • Nitrite
  • Oxalate
  • Phosphate
  • Thiocyanate

Ammonium ion

It is a positively charged polyatomic ion with the chemical formula NH4+. It has a tetrahedral geometry(sp3 hybridization).It is formed by the reaction of aqueous ammonia and acids.

NH3 + H+ → NH4+

If we add concentrated hydroxide to an aqueous ammonium solution, then red litmus paper(moistened) turns blue; this is used to detect ammonium ions. It is an important nitrogen source for many plant species. The lewis structure of ammonium ion is given below:

polyatomic ion examples
LEWIS STRUCTURE OF AMMONIUM ION

Acetate ion

It is a negatively charged ion with the chemical formula CH3COO–. It is formed by the removal of a proton from the carboxy group of acetic acid.

To detect acetate ions, we add ethanol and conc. H2SO4 to a salt of acetic acid and heat it, the presence of fruity odor confirms the presence of acetate ions. It is used to line fabrics in robes and loungewear. It is also used as a precursor of acetyl-CoA for synthesizing fatty cells. The lewis structure of acetate ions is given below:

Acetate ion
LEWIS STRUCTURE OF ACETATE ION

Carbonate ion

It is a negatively charged ion with one chemical formula CO32- arranged in trigonal planar geometry.

Carbonate ions soften water and are used in the production of paper and glass industry. When we add dilute sulfuric acid to a salt containing carbonate ion, there is an effervescence with the evolution of carbon dioxide (a colorless and odorless gas). The lewis structures of carbonate ion are shown below:

Carbonate
LEWIS STRUCTURES OF CARBONATE ION

Chromate ion

These are the anions present in chromate salts. Its chemical formula is CrO42- and is a powerful oxidising agent and is bright yellow in colour.

They have tetrahedral geometry. Chromates are used in inks and dyes as pigments, in chrome plating (to get rid of corrosion), leather finishing and metal rust proofing. Pb2+, Ag+ and Ba2+ forms sparingly soluble precipitate with chromate ions. PbCrO4 and BaCrO4 form a yellow precipitate while Ag2CrO4 forms a brick red precipitate. This is used to detect chromate ions.

Chromate
LEWIS STRUCTURE OF CHROMATE ION

Cyanide ion

It is an anion containing the C≡N group (carbon triple-bonded to nitrogen) and is conjugate base of hydrogen cyanide. It has a linear geometry with a negative charge on the carbon atom.

The major use of cyanide is in mining gold and silver. Cyanide ions are also used to produce CN-containing compounds (usually nitriles). We add conc. H2SO4 to cyanide-containing solution; hydrocyanic gas is formed and the reaction color changes from pale green to blue.

Cyanide1
LEWIS STRUCTURE OF CYANIDE ION

Hydroxide ion

It is an anion with the chemical formula OH– derived from base (NaOH, KOH, Ca(OH)2) or water by loss of a proton.

Hydroxides are used in manufacturing paper, pulp, soaps and detergents. If we add an ammonium salt to a base containing hydroxide ions, ammonia gas will be evolved (which can be detected by bringing a glass rod dipped in concentrated HCl, ammonia gas will give solid white fumes).

HYDROXIDE
LEWIS STRUCTURE OF HYDROXIDE ION

Nitrate ion

It is an anion with the chemical formula NO32- , formed by the loss of a proton from nitric acid.

Nitrates are used as oxidizing agents, fertilizers, and explosives. To detect nitrate ions, we add concentrated H2SO4 to a salt containing nitrate ions, cool it, and then add FeSO4 to it. A dark brown ring confirms the presence of nitrate ions.

NITRATE1
LEWIS STRUCTURES OF NITRATE ION

Nitrite ion

Nitrite is an anion with the chemical formula NO2– and is derived from nitrous acid. It has a bent shape with sp2 hybridization.

They are used to prepare azo dyes and other colorants. When dilute H2SO4 is added to a salt containing nitrite, reddish-brown fumes of NO2 gas are observed.

NITRITE
LEWIS STRUCTURE OF NITRITE ION

Oxalate ion

It is an anion of a dicarboxylic acid with the formula C2O42- which is used as a human and plant metabolite.

They are used as reducing agents (strontium and barium oxalate), in photography and to remove ink stains. The oxalate anion exists in planar geometry (like potassium oxalate) and non-planar in other cases (like caesium oxalate). Free oxalate anion has an orthogonal D2h structure.

The addition of concentrated H2SO4 to an oxalate salt produces colorless, odorless gas (CO2), which turns lime water milky (detection of oxalate ion).

LEWIS STRUCTURES OF OXALATE ION

Phosphate ion

It is an anion with the chemical formula PO43- and is derived by the loss of three protons from phosphoric acid. It has a tetrahedral geometry.

It is used in fertilizers. Phosphate ions provide energy to cells and are important for bone and teeth formation.

For detection of phosphate ions, we add conc. HNO3 and ammonium molybdate to the solution containing phosphate ions and heat it. A canary yellow precipitate of ammonium-phosphomolybdate is observed, which confirms the presence of phosphate ions.

PHOSPHATE1
LEWIS STRUCTURES OF PHOSPHATE ION

Sulfate ion

It is a sulfur oxoanion with the chemical formula SO42- and is derived from sulphuric acid by deprotonation of both OH groups. It has a tetrahedral geometry.

It has various industrial uses. It is used to produce plaster, as an electrolyte in galvanic cells, and as a detergent in shampoo formulations. An aqueous solution of salt-containing sulfate ions(acidified with acetic acid) reacts with BaCl2 to form a white precipitate of BaSO4.

SULFATE
LEWIS STRUCTURES OF SULFATE ION

Sulfite ion

This is also a sulfur oxoanion but with the chemical formula SO32- and is a conjugate base of bisulfite.

They are used as preservatives, bleaching agents and dechlorinating agents. It has trigonal pyramidal geometry. When a solution containing sulfite ions reacts with dilute H2SO4, SO2 gas evolves, which has a smell of burning sulfur. This gas turns potassium dichromate paper(acidified with dil. H2SO4) green confirming  the presence of sulfite ions.

Sulfite
LEWIS STRUCTURES OF SULFITE ION

Thiocyanate ion

It is an anion with the chemical formula SCN– where the negative charge is shared by nitrogen and sulfur equally(ambidentate ligand). It is derived from thiocyanic acid.

It is utilized in bleach and disinfectants and to prepare silver thiosulphate, which inhibits corrosion in steel. It is an important spectrophotometric reagent. A blood-red color is observed when iron nitrate is added to a solution containing thiocyanate ions.

Thiocyanate
LEWIS STRUCTURE OF THIOCYANATE ION

Conclusion

We discussed thirteen important polyatomic ion examples, their uses, molecular geometry, lewis structure and qualitative tests for their detection.

AlCl4- Lewis Structure, Geometry: 9 facts you should know.

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Tetrachloroaluminate anion 2D A

In this article we are going to see about alcl4- lewis structure and some important facts around this.

Lewis structure or lewis dot structure is a simple representation of the electronic structure of a molecule that briefs about the number of bonds formed, number of bond pairs required to fulfill the octet rule and lone pairs available. This method of drawing a molecule helps in simple representation of a molecule by allowing to guess its structure or shape.

 An atom always tend to arrange 8 electrons around themselves to acquire a stable or a noble gas configuration with some exceptions like when a molecule is electron deficient; when it has odd number of electrons; or molecules that has extra electrons in their valence shells. E.g., BH3 ,SF6 ,NO etc.

Methods to draw a Lewis structure : 

  • First count the total number of valence shell electrons available for each atom.
  • Choose the least electronegative atom as the central atom and draw the remaining atoms around the central atom and start by forming a covalent bond (a bond requires two electrons). An atom will always try to fulfill its octet or expand its octet if required. 
  • The remaining electrons not forming covalent bond will stay as lone pair of electrons.

Note: Elements having expanded valence shells like 3d elements, it can exceed the octet rule like SF6 , PF5  or elements with fewer valence electrons can have incomplete octet like H2 .

alcl4- lewis structure
AlCl4- Lewis Structure from Wikipedia

Lewis structure of AlCl4– :

Aluminium belongs to 13th group and 3rd  period. It has 3 valence electrons, with an empty 3d shell that can expand its octet if required ( but it does not due to steric hindrance.) It has 4 Cl atoms with electronegativity 3.16 and Al with electronegativity 1.61, therefore, choose Al as the central atom. It has a total of 32 valence electrons from 4 Cl atoms, 1 Al atom and a negative charge. Draw a covalent bond between each atoms with the central atom to fulfill the octet. In doing so, we get 12 lone pairs of electrons, 3 each on the surrounding atoms. Each atom’s octet is filled thereby not violating the octet rule.

alcl4 1

Formal Charge of AlCl4– :

It briefs about the hypothetical charge acquired by an atom in a molecule if the electron pairs were shared evenly between the atoms to fill its valency completely.

Generally, formal charge can be calculated mathematically by the formula :

Formal charge = (Number of valence electrons in a free atom of the element) – (Number of unshared electrons on the atom) – (Number of bonds to the atom)

In addition, Charge on the molecule= sum of all the formal charges .

Formal charge of Al atom = 3-0-4 = -1

Formal charge of Cl atom = 7-6-1= 0

Since all the chlorine atoms are equivalent, hence we assign 0 formal charge to all the Cl atoms ( meaning they are neutral )    

alcl4 2

Resonance structure of AlCl4– :

Aluminium has valence shell configuration 3s2 3p1 . It generally shows a covalency of 3 acting as a Lewis acid but it can accommodate extra electrons in its empty 3d orbital and two 3p subshells thereby extending its covalency to 8. However, Al has small size which cause steric hindrance with greater covalency.

Therefore, it sticks to its minimum covalency of 3 and if an extra electron is provided by a donor, it can extend its covalency upto 6. In our case of Tetrachloroaluminate, it extends its covalency to 4. Therefore, it doesn’t form any double or triple bonds due to steric hindrance. All the 4 resonance structures are equivalent which is shown below :

  This form is the most stable orientation of AlCl4– . The `arrow’ is the coordinate/dative bond.

  Note: Dative or coordinate bonds are formed by sharing two electrons covalently by an atom to the nearest neighbouring atom. 

alcl4 3

AlCl4–  valence electrons, bond pairs and lone pairs of electrons :

It has a total of 32 valence electrons. 

It has 4 bond pairs ( that is formed by a single covalent bond here) and 12  total lone pairs of electrons ( that did not participate in bond formation). 

Note: In actual sense of chemistry, there is a coordinate bond formation between AlCl3 and Cl– ( electron pairs shared completely by chloride ion, a type of covalent bond).

Hybridisation, Shape and Angle of AlCl4– :

There is a simple rule or equation to be followed to find out the hybridization of molecules real quick. 

Hybridisation of a molecule = ( Valence electrons of the central atom + Number of monovalent atoms attached to the central atom + Negative charge on the molecule – Positive charge on the molecule )/2

Here, AlCl4– hybridization = (3+4+1)/2 = 4 i.e., sp3

It is a tetrahedral structure and not square planar as alkaline earth metals do not necessarily exhibit square planar complexes formation since they do not contribute to the overall CFSE value.

It has sp3 hybridisation with tetrahedral geometry and bond angle of 109’5” . It has 4 single bonds with bond length of Al-Cl at around 2.06-2.08 armstrong.

alcl4 5
Molecule Tetrachloroaluminate, AlCl4–
Type Polyatomic ion
Hybridisation sp3
Shape Tetrahedral
Bond angle 109’5”
Bond and Lone Pairs 4 , 12

Is AlCl4–  stable ? 

Yes, AlCl4–   ion is found to be stable. It has high thermal stability, high conductivity, low melting point which make it an excellent choice for further study and investigation. It has its octet complete making it a stable compound.

Note : Formation of AlCl4– –  As AlCl3 has empty p orbital that can accept another electron to complete its octet, it reacts with a chloride ion easily to form a more stable ionic molecule. They are also observed as an intermediate species during many organic reactions like Friedel Crafts alkylation, acylation etc.

Is AlCl4–  ionic or covalent ?

Tetrachloroaluminate is a covalent molecule. It is a polyatomic anion where a chloride ion shares two electrons by a covalently coordinate bond to AlCl3(covalent) . 

Note: A polyatomic anion with a metal cation yields an ionic compound.

Uses of AlCl4– :

  1. Due to its thermal stability, low melting point and low vapor pressure, it is widely used as an ionic electrolyte for redox reactions and chemical reactions.
  2. It is used in solvent extraction, organic catalysis reaction, in Al dual-ion batteries.
  3.  They are used in batteries

Conclusion :

AlCl4– or Tetrachloroaluminate is a sp3 hybridised polyatomic molecule with tetrahedral geometry containing 4 bond pairs and 12 lone pairs which is widely used for commercial and industrial purposes.

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