Hi...I am Triyasha Mondal, pursuing M.Sc in Chemistry. I am an enthusiastic learner. My specialization is in physical chemistry.
Let's connect through LinkedIn:
Iron (III) Bromide (FeBr3) has quite stable polymeric structure, generally known as Ferric bromide.
The brown color solid inorganic compound can be used as oxidant, removes electron in redox reaction or used as catalyst in bromination where concentration of FeBr3 remains same after and before reaction but it affect the rate of reaction.
How to draw febr3 lewis structure?
With using the atomic symbols, the FeBr3 lewis structure shows the distribution of the outer orbit electrons in formation of the molecule, as an ionic compound for the electronic transition the charges are shown upper outside of bracket around atom.
Count the valence electrons of the molecule:
In modern Periodic table the halogen atom Bromine is in group 17 with outer shell electronic distribution in 4s, 4p orbital: [Ar] 3d10 4s2 4p5 where Iron is in group no 8 and being transition metal with incomplete electronic arrangement in 3d orbital; [Ar] 3d6 4s2.
Outer orbit electrons are mostly loosely bounded by nucleus force but in addition the d orbital being more diffused in nature it can easily exchange electron cloud. Orbital is energetically most stabilized in their full filled or half filled electronic configuration.
Cover the Octet rule:
To form energetically stable molecule (unreactive), all atoms must have full filled last orbit with eight electrons by transferring electrons, referred as Octet rule.
Iron transfers three electrons from 4s and 3d jointly results remaining of five electrons in 3d orbital, which is a stabilize configuration for the half filled 3d orbital and forms Iron (III). The electrons are accepted by the vacant 4p orbital to satisfy the stabilization rule.
Calculate the formal charge of the individual atom:
As FeBr3 is ionic molecule so the formal charge is not calculated where we consider the electro-negativity of individual atoms, opposite of the covalent molecule, where the electron pulling ability of atoms is consider equally.
Skeletal formation of the BrCl lewis structure:
FeBr3 is a tetra-atomic molecule, where Iron (Fe) is the central atom, more electro positive than halogen atom. The atomic symbols are written inside brackets, by electron transferring they possess charges which are shown outside of bracket.
When metal and non-metal elements from molecule, there fully electronic transfer is occurred from metallic orbital to vacant non-metallic orbital in spite of electron cloud sharing, produces ionic compound.
Bromine is a non metal element where Iron is a transition metal element, are joined together by fully transferring electron cloud of Fe having diffused 3d orbital electron to vacant 4p orbital of halogen atom, hence FeBr3 lewis structure is ionic in nature.
Being an ionic molecule we can see the melting point of the FeBr3 lewis structure is quite high, (200 ̊ C) as the opposite charges are bound by strong electrostatic force and are arranged in a proper design of repetition in crystalline structure, result in a high lattice energy.
Is febr3 an acid or base?
FeBr3 lewis structure contains Fe atom which is a transition metal having 3d orbital, and Fe combines with electro-negative halogen atom Br, makes the molecule a lewis acid, capable of accepting electron.
In metal halide FeBr3, because of difference in electron cloud pulling capability Fe contains partial positive charge and because of having vacant 3d orbital it adopted electron pair accepting ability, as a result used as a catalyst in aromatic bromination reaction.
In bromination of Benzene FeBr3 is used with Br2 for its strong lewis acid nature, it polarizes the Bromine molecule with partial charges and combine with it by accepting electron cloud which makes easy for benzene electron cloud to do nucleophilic attack.
FeBr3 solubility in any solvent is measured by the amount of concentration that can be dissolved until the equilibrium occurs means the concentration of reactant compound and the product compound result balance in that condition (like temp.)
Solubility in water:
As a polar solvent, water (H2O) has partial charges over it for difference of electro-negativity between atoms and can dissolve polar molecules. Because of greater power of pulling electron cloud toward it, Oxygen carries partial negative charge.
FeBr3 has Iron atom which has (+3) charge over it as it transfer two 4s and one 3d electrons and three Bromine atoms which accepts electrons so that have negative charge (-1), so charge separation is there and as a result can easily dissolve in water like polar solvent.
Solubility in ether:
Ether has polar Carbon- Oxygen bond (Oxygen is highly electronegative than Carbon) as its molecular formula is R-O-R´ where R and R´ are two different alkyl groups, so it has some dipole moment causes ether a polar solvent.
FeBr3 is an ionic compound, contains oppositely charged ions bounded with electrostatic force, so can dissolve in polar solvent. But the polarity differs with length of hydrocarbon chain of ether as with increasing it the interaction between ether and FeBr3 decrease.
Solubility in ethanol:
Methanol, ethanol or propanol have small hydrocarbon chain (hydrophobic part) with the hydroxyl group (-OH) which is polar for Oxygen atom with high electro-negativity (electron pulling ability) result in becoming polar solvent.
So ionic compound like FeBr3, having charge distribution can soluble in ethanol.
Is febr3 a salt?
A salt have two portion called cation with positive charge and anion with negative charge, the amount of the charge will be equal so that the salt molecule will be neutral without any charge. If they contain the ionic part forming with one atom, are called monatonic salt.
In neutralization reaction, FeBr3 is formed having oppositely charged metal ion iron (III) and non-metal bromide ion, and as one iron atom combines with three bromide ions, produces neutral salt. As the ionic parts are not combination of atoms, the salt is monatomic.
In the neutralization reaction of strong acid (HBr) and weak base (Fe(OH)3) where the strong acid is neutralized and forms the ionic compound as salt FeBr3 and water.
Being an ionic compound Salts have some crystalline structure where the ions are arranged in certain manner of repetition which also can be seen in case of FeBr3 lewis structure where it crystallizes in the trigonal R-3 space group.
Conclusion:
Having transition metal, iron with ‘d’ orbital effects properties of FeBr3.
Uric acid is a well known chemical compound of mammal body, which appears as white crystals physically.
For the heterocyclic formation with two configuration, keto and enol, we see many properties which are mentioned in this article.
Skeletal of Uric acid:
Two cyclic rings are connected to produce the Uric acid structure; among these one is six member ring and another is five member ring with the chemical formula C5H4N4O3.
As the ring contains atoms other than Carbon and Hydrogen which are Nitrogen and Oxygen, so the compound Uric acid is called hetero- cyclic compound.
The IUPAC (International Union of Pure and Applied Chemistry) name of Uric acid is 7,9-Dihydro-1H-purine-2,6,8(3H)-trione and the other chemical names of this structure are 1H-Purine-2,6,8-trione, 2,6,8-Trioxypurine, 2,6,8-Trioxopurine etc.
uric acid structure
Formation of Uric acid:
Uric acid is produced through a complex pathway from purine nucleotides of animal proteins mainly in liver, intestines where different types of enzymes are involved in the process like deaminase, xanthine oxiddase, catalysis the reaction.
Uric acid is synthesised mainly from two substances, Adenosine monophosphate and Guanine monophosphate which are converted to Adenosine and Guanosine respectively by nucleotidase.
In the further reaction inosine and guanosine produce hypoxanthine and guanine respectively; from where Hypoxanthine is oxidized (removal of electron) by the enzyme called xanthine-oxidase to Xanthine and Guanine is deaminated to Xanthine.
Finally Xanthine is oxidized means removal of electrons or addition of Oxygen, by the enzyme, Xanthine oxidase to form the product Uric acid.
Ions and salt of Uric acid:
Uric acid is a weak acid which can release Hydrogen ion (H+) and produced its ion Urate with having negative charge on it, hence both form weak acid-conjugate base pair. The negative charge can delocalize through ring, give stability to structure.
In blood Sodium ion (Na+) is one of the main cation present which can interact with the Urate ion as it is negatively charged compound and both the oppositely charged particles form the salt called monosodium urate.
ion of uric acid structure (urate ion)
Tautomerism of Uric acid structure:
If structural isomers (structures with same chemical formula but different bonding patterns) rapidly inter converted then they called tautomers and the chemical process is called tautomerization.
In Uric acid structure, there are three carbonyl carbon atoms (keto form) which can convert to enol form by taking Hydrogen ion from the hetero atom Nitrogen. Hence the chemical formula remains same, we see two different configurations.
If the heterocyclic compounds have conjugated planar ring system means the pi (π)-electrons can spread over the ring with (4n + 2) pi (π) electrons the hetero cyclic compounds become aromatic in nature.
In Uric acid there are two cyclic rings present and also we see two tautomeric form of Uric acid, one is Lactam and another is Lactim. In lactim structure we see the extended electronic conjugation means they can delocalize.
Solubility of Uric acid:
Solubility of Uric acid and its salts depend on the concentration that soluble in the solution at standard condition, for water which is very low. They can be soluble in hot water rather than cold water.
This low solubility can cause easy crystallization for the salts of the Uric acid. In Ethanol the solubility of Uric acid and its salts are nearly negligible.
Acidic nature of Uric acid:
Acidic nature of the Uric acid depend on the ability to release the Hydrogen ion (H+) in medium as the measurement parameter of acidity is pH which is the negative logarithm of the Hydrogen ion (H+) concentration.
If any compound release large concentration of (H+) ion the pH value will be small, called strong acid where as if the opposite happen, the compound becomes weak acid. The Lactim form of Uric acid have three (-OH) group, one in five member ring, others in six member ring.
When Urate ion forms the (-OH) groups of the six member ring don’t release any (H+) ion, only the (-OH) group of the five member ring release (H+) ion due to stabilization factor.
Hence for low hydrogen ion concentration Uric acid becomes a weak acid with pKa around 5.6. However depending on the medium uric acid can release more than one proton (H+) and become polyprotic acid.
Effect of excessive Uric acid in blood:
Uric acid is a sparingly soluble molecule, hence if its concentration in blood increases, due to the low solubility it start to form sharp crystals as precipitation around the joints, causes pain.
Hence taking of Purine rich foods in large amounts daily must be avoided. Low level of Uric acid in blood also causes different problems.
Conclusion
For the tautomeric forms of uric acid it shows different properties.
Sialic acids are found in animal tissues and other microorganisms, and human brain cells contain that in high amount.
For the different linkage, different verities of Sialic acid structure can be seen which are part of glycoproteins, glycolipids, produces different biochemical functions which are described in this article.
Structure of Sialic acid:
Sialic acid structure is an alpha- keto acid sugar with a skeletal of nine Carbon atoms, which is obtained from a molecule, neuraminic acid by replacing with amino group of one of its hydroxyl group (-OH).
Nearly 50 types of derivatives of Sialic acid structure are known now. In the heterocyclic ring, at the position of Carbon-5, amino group is present and at Carbon-1 a carboxyl group is present which contains negative charge, makes Sialic acid strong organic acid.
The amino group is usually acetylated means replacing one Hydrogen atom with acetyl group, which produces N-acetylneuraminic acid, the common form of Sialic acid.
Sialic acid structure
Numbering of the Carbon atom of Sialic acid structure:
The numbering of the Carbon atoms of the Sialic acid structure starts from the Carboxylate Carbon atom means it is C-1 and then continues along with the chain.
There are two structural configurations, one is alpha (α) anomer and another is beta (β) anomer. If the Carboxylate group at C-1 is in the axial position then the configuration is called σ-anomer and if it is on equatorial position then called β-anomer.
alpha configuration of sialic acid structure
Mainly over 90% the beta anomer is formed.
beta configuration of sialic acid structure
Diversity in the linkage:
Different types of alpha (α)- linkages are formed Carbon-2 of Sialic acid and sugers, beside this common linkages are occurs to the Carbon-3 or Carbon-6 positions of the galactose residues.
Sialic acids can cover the internal positions within glycans, among which most common is Sialic acid is when attached to another, usually at the Carbon-8 position. The Carbon-5 position can be attached with N-acetyl group or a hydroxyl group.
The Carbon-1 which is a carboxyl group generally remains as ionized form at physiologicalpH, but it also can be condensed in the lactone form which has hydroxyl group or in the lactam form with a free amino group at Carbon-5 position.
Chemical diversity of Sialic acid structures contribute in the formation of the variety of glycan structures on the cell surface, results in the formation of different cell types.
Biosynthesis of Sialic acid:
Sialic acid is synthesized from glucosamine 6 phosohate and acetyl coenzyme A (deliver acetyl group) through a transferase, enzyme which catalyses the transfer of a specific functional group; results N-acetylglucosamine- 6-P formation.
This product becomes N-acetylmannosamine-6-P through epimerization process, where epimer is transformed to its diastereomeric other part. This reacts with phosphoenolpyruvate (ester derivative from enol of pyruvate and phosphate).
The above reaction produce N-acetylneuraminic-9-P compound which is called Sialic acid. To turn it active for biosynthesis process a monophosphate nucleoside is added with the Sialic acid, results of the formation of cytidine monophosphate-sialic acid.
Function of Sialic acid:
The functions of Sialic acids are generated due to the strong electro- negative charge over it, as it can bind and transport ions (cation) through electrostatic attraction and also stabilize the conformation of proteins including enzymes also.
For the charge Sialic acid carrying, it can increase the viscosity (the resistance power to deformation) of mucin, high molecular weight proteins that can form gel (main component of mucus), protects molecules and cells from attack by glycosidases, increases lifetime.
It controls the affinity of the receptors, protein structures which can bind with chemical messengers and cause change in electrical activity of cell and can report to the modulate processes involve membrane signaling, growth etc.
Hydrophilic nature of Sialic acid structure:
On the surface membranes glycoconjugates are found, containing Sialic acid rich oligosaccharides which help to keep water molecule at the surface of the cell.
As the carboxylate group of the Carbon-1 remains as negatively charged and the Sialic acid remains in the terminal position of carbohydrate chain, so the Sialic acid rich regions create a negative charge on the cell surface.
Water is polar molecule which has partial positive charge on the Hydrogen atoms, so it is attracted to the negatively charged cell surface and membrane.
Acts as masking component:
Erythrocytes and other blood cells are covered by a dense layer of Sialic acid during their normal life time and is removed step wise by hydrolysis. Finally the unmasked cell is destroyed by phagpcytosis, so Sialic acid protects them.
Acts as recognition sites:
Sialic acid structures (components of glycan ligands) are recognized by molecules which belong to Lectin.
Plants don’t have Sialic acids, so these Lectins can help to defense against the microorganisms which contain Sialic acids. In Sialic acid structure the Carbon-1 which contains Carboxylate group (negatively charged) is proven critical for recognition.
Conclusion
For the carboxylate group present in the Sialic acid structure, the negative charge density generated which is one of the cause of the different functions.
BrCl is a strong oxidizing agent mainly used for analytical chemistry purpose.
Bromine Chloride is highly reactive species, in fact it decomposes at room temperature. We can explain the properties of it with the help of the BrCl lewis structure.
How to draw Brcl lewis structure?
With using the atomic symbols of Chlorine and Bromine, the BrCl lewis structure shows the distribution of the outer orbit electrons, spread around the particular atom in the molecule. To draw the structure some steps are followed;
Count the valence electrons of the molecule:
In modern Periodic table the halogen atoms both Chlorine and Bromine are in group 17. The outer shell electronic distribution of Chlorine atom is [Ne] 2s2 2p5. Bromine atom contains its outer orbit electrons in 4s and 4p with the arrangement: [Ar] 3d10 4s2 4p5.
Hence Chlorine and Bromine jointly have fourteen loosely bounded electrons means seven pairs of electrons for molecule formation.
Skeletal formation of the BrCl lewis structure:
BrCl is a diatomic molecule, so here no need to choose the central atom as for polyatomic molecule. The atomic symbols of the halogen atoms written sidewise, for Chlorine ‘Cl’ and Bromine ‘Br’ and the valence electrons are arranged accordingly.
To form energetically stable molecule, both the halogen atoms must have to cover the last orbit with eight electrons, so they share electrons cloud. Hence a single bond (sigma bond) forms, that is drawn by a single line in BrCl lewis structure.
Calculate the formal charge of the individual atom:
After making a lewis structure we have to calculate the formal charge of the halogen atoms. If both Chlorine and Bromine have zero formal charge over them, then only the BrCl lewis structure will be a stable one.
Brcl lewis structure resonance
When an atom has unshared electrons in a molecule, these can be able to spread into vacant orbital of neighbor atom of the molecule and we get canonical resonate structure, which are stabilized by conjugation energy.
Both electronegative halogen atoms form the BrCl lewis structure through the sigma bond formation. If the electron cloud delocalizes there will be electronic repulsion which destabilize the structure. So we get no resonate structure.
As AXE3 type molecule (where A is Br, X is Cl and E represent the lone pair) the geometry of the molecule should be tetrahedral but both of the halogen atoms contain unshared negative charged cloud. For the equal repulsion we get a linear shaped molecule.
Considering the bonding electron cloud is equally distributed, formal charge can be calculated by the equation, f = (n – lp – (bp/2)), where ‘n’ is outer shell electrons, lp is nonbonding electrons and bp is electron cloud distributed for bond formation.
From the Periodic table, we know that halogen atoms, here both Chlorine and Bromine have seven electrons in their last filled electronic shell. In the BrCl lewis structure, after the bond formation each halogen atom has three pairs of electrons which remain as unshared electrons.
Hence for both Chlorine and Bromine atom, formal charge in above BrCl electron dot structure is, {7- 6 – (1/2×2)} = 0.
As the atoms of this above electron dot structure have “zero” (0) formal charge, hence this skeletal of BrCl lewis structure is very much energetically stable compared to the other possible electron dot structure.
Brcl lewis structure angle
Bond angle is measured as the average angle between the adjacent covalent bonding orbital which changes with different configurations of electron dot structure.
Being a diatomic molecule, the two atom’s 3p and 4p orbital respectively are overlapped directly, results a sigma bond formation with sharing the electron cloud. Hence in 3D space they belong to a plane, where the angle between them can only be 180 degree.
In main group elements, eight electrons are arranged in their valence orbit, which describes electron’s wavy nature, in such a way that the molecule becomes inactive and become stabilized which is termed as Octet rule.
In the BrCl lewis structure both Bromine (Br) and Chlorine (Cl) each have loosely bounded seven negative particles in their outer energy shell. By sharing the electron cloud they easily cover up the outer orbit and become energetically stable like inert gas.
Brcl lewis structure lone pairs
Electron cloud that are loosely bounded by nuclear force, participate to reaction for bond formation but some remain unshared during the overlapping. These electron cloud are called lone pairs with higher electronic density.
To satisfy the Octet rule for stabilization, the two atoms in BrCl lewis structure need one electron each. Hence among the participated outer filled loose electrons one pair is utilized in bonding and remaining six pairs of electrons concentrate over their parent atom.
Brcl valence electrons
The outer shell electrons which are loosely bounded by nuclear force of attraction due to the shielding effect of the inner shell electrons, called valence electrons. They can absorb energy to be in excited state to participate in chemical reactions.
In BrCl lewis structure the two halogen atoms contain fourteen electrons, where Bromine atom has seven negatively charged particles in the 4p orbital whereas the Chlorine atom has seven electrons in the “p” orbital of 3rd orbit.
The core orbital electrons of the halogen atoms are tightly bounded to their nucleus by strong nuclear force, but with increasing the number of orbits (electronic energy level) from the nucleus, these force lowers.
For this reason the outer shell 3p and 4p electrons of Chlorine and Bromine atoms are loosely tied up. They act as free energy particle and can move to form bond between the atoms.
Brcl hybridization
Hybridization is the concept in Valence Bond Theory, for covalent compounds where atomic orbitals intermixes to form energetically equivalent orbital, so that it become easy for overlapping during molecular orbital formation.
To form Brcl lewis structure sigma bond the energetically comparable ‘p’ orbitals directly overlapped, where only one electron is needed per atom and they have vacant space. Hence no requirement for hybrid orbital formation.
Brcl solubility
BrCl solubility in any solvent is measured by the concentration which can be dissolved until the equilibrium occurs in the solution. This depends on facts like the type of bonding between the inter-halogen compound, polarity, temperature etc.
Having the electro-negativity difference, polarity exists in the molecule results salvation in polar solvent water with the amount of 8.5 g/L. It is also soluble in Carbon disulfide (CS2) and ether type solvents.
Is Brcl ionic?
If atoms of a molecule are bonded together by electrostatic force, with completely electronic transformation, then it is ionic compound, can be shown between metal and nonmetal elements. In spite this if electron cloud is shared between atoms, is called covalent compound.
Bromine and Chlorine both are non metal elements, are joined together by sharing electron cloud of their outer most orbital 4p and 3p respectively. So the electron cloud of the bond is not concentrate only over particular atom. It is a covalent molecule, not ionic.
Being a covalent molecule we can see the melt point and boiling point of the BrCl lewis structure is very less, (-54 ̊ C) and 5 ̊ C respectively where these properties of an ionic molecule is very high as the opposite charges are bound by strong electrostatic force.
Is Brcl acidic or basic?
BrCl is a very reactive inorganic species, which decomposes even at room temperature.
Its solubility in water is good (8.5 g/L) which produces hypobromous acid. So the water solution of BrCl is acidic in nature.
Is Brcl polar or nonpolar?
In the BrCl lewis structure, electro-negativity of Chlorine (Cl) is 3.16 and that of Bromine is 2.96 on the basis of Pauling scale, so difference creates charge separation in bond and which leads the molecule to be polar.
In the molecule both the atoms are electron-egative in nature as they belong to the group no 17 ( right column of the Periodic table), but for having ‘d’ orbital the outer orbit electrons of Bromine are more diffused, hence Chlorine has more tendency to pull the bonding electron cloud.
BrCl lewis structure is linear, so the vector direction of electro-negativity towards the Chlorine atom can’t be cancel out which results of dipole moment as it is measured by the multiplication of the charge separation (δ) and the distance between the charges (r) and BrCl become polar.
The diatomic inter-halogen compound BrCl is unstable at room temperature, and the other properties are easily explained by the electron dot configuration.
CaBr2 is mainly used in high concentration of water solution for drilling purpose or as ingredient of food preservative or for prepare freezing mixture.
Calcium Bromide is a white powder with crystalline texture and having other properties. We can explain these properties with the help of the CaBr2 lewis structure and the behavior of the outer orbit electrons which are shown in the structure.
How to draw Cabr2 lewis structure?
CaBr2 lewis structure is a presentation way for showing the bonding pattern between the atoms with the loosely bonded outer shell electrons. Calcium is in group 2 in the Periodic table with the electronic arrangement, [Ar] 4s2.
Bromine is a non-metal element with the arrangement of electrons, [Ar] 4s2 3d10 4p5. So the molecule contains total sixteen loosely attached electrons around the particular nucleus. Having less capability to pull the electron cloud, Calcium becomes central atom.
Bromine atoms remain to the opposite sides of Calcium atom and transferring electron cloud form ionic bonds. The sigma bond between the Calcium and Bromine atom is shown by a single line in the CaBr2 lewis structure and the ions are written inside bracket.
Resonance is a concept of VBT where the electron cloud can delocalize to the neighbor vacant orbital and form a canonical structure which contributes in the main structure of the molecule.
In an ionic molecule like CaBr2 where the electron cloud is fully transferred to the nonmetal atom vacant shell, there is no chance to delocalize the electron cloud again to form resonating structure.
Cabr2 lewis structure shape
Molecular geometry and shape are two slightly different things where which we can clearly understand by the hybridization of central atom of a covalent molecule. Lone pair distorted the geometry by steric repulsion and introduces the molecular shape.
As an ionic compound the oppositely charged ions bounded by strong force in the lattice structure. So individually we can’t define the shape of the ionic molecule by hybridization method. Small unit of molecule repeatedly arranged to form the lattice structure.
But it can be said that the molecule has two ionic where Calcium atom is in the central position.
In the assumptions of the formal charge calculation of particular atom of a molecule, it is stated that bonding electron cloud must be distributed equally over the two atoms which only possible in the case of molecular orbital overlapping.
In CaBr2 lewis structure electron cloud is fully transferred so two oppositely charged ions, i.e, Cation and Anion produced. So here we count the oxidation number in state of the formal charge, where oxidation number is the charge gained by an atom during electron transfer.
Calcium losses its two electrons, so electronic charge decreases in the atom, results in a positively charged ion formation with charge (+2), which means oxidation number of Calcium atom in the CaBr2 lewis structure is (+2).
We can calculate the oxidation number of Bromine atom in the CaBr2 lewis structure by the same process which will be (-1). Hence overall the molecule CaBr2 is electronically nutral.
Cabr2 lewis structure angle
Bond angle is a parameter of a molecular structure which is covalent in nature. This covalent type of bonding is formed by hybrid orbital overlapping, so mainly electron cloud is shared between the adjacent atoms, which are not possible in ionic compound.
In ionic compound the oppositely charged particles tightly packed by strong attracting force and arranged in a particular geometric way to produce lattice. So no such angle is measured in this type of ionic compound.
Cabr2 lewis structure octet rule
Eight electrons are arranged in the orbit which is wave functions of specific energy in such a way that the molecule gets stabilization energy. This termed as Octet rule. Nobel gases have this type of energetic stability.
Calcium have two extra electron in 4s orbital than its nearest inert gas Argon. These are transferred to two Bromine atoms which lack one electron to gain this stability. By this method both atoms bring stability in CaBr2 lewis structure.
Cabr2 lewis structure lone pairs
After molecular formation some atom contains such valence electron pair which only spread over that atom and don’t take part in the bond formation, are known as lone pair electron. These electron cloud concentration is higher than bonding electron.
Lone pair electron has negative character so can posses in a particular vector direction. In CaBr2 lewis structure two electrons of outer shell 4s orbital of Ca are fully transfer to vacant 4p orbital of Br and it become ion with positive charge so there is no lone pair over Ca atom.
As lone pair electron cloud spread over more space with higher charge density, it can disturb the bond angle as well as the shape of the molecule.
Cabr2 valence electrons
The last shell containing electrons which are energetically approachable for excitation in any chemical reaction are known as valence electrons. This are far apart from the nucleus and for the shielding of inner orbital electrons, are loosely bounded.
Being a group 2 element in 4th Period, Calcium has two loosely bounded electrons in 4s orbital whereas Bromine contains seven that type of electrons in 4s and 4p orbitals. So total sixteen electrons loosely bounded by the nuclear force are present in the molecule.
Among these valence electrons two pairs are joined in the two Ca-Br bonding. Other remains equally distributed over two Br atoms as non-bonding electron pairs.
Cabr2 hybridization
Hybridization is the concept for covalent compounds where energetically equal but not equivalent atomic orbitals intermixes to form equivalent orbital so that molecular orbital can form by greater atomic hybrid orbital overlapping.
Calcium Bromide is ionic compound where the atoms are connected by totally transfer of electrons from metal atom, Calcium 4s orbital to the halogen, Bromine atom p orbital. Hence here is no requirement for formation of any hybrid orbital for bonding generation.
Cabr2 solubility
Solubility in water depends on the fact that if the solute contains any polar part in it as water itself is a polar molecule. Being an ionic compound CaBr2 can attract water molecules and form hydrated structure generally with six water molecules.
It also can soluble in acetone solution.
Is Cabr2 acidic or basic?
If we notice the formation of the CaBr2 lewis structure, we see that it forms by the chemical nutralization reaction between a strong acid (high capability of accepting electron) and strong base.
Hence CaBr2 is nor acidic or basic, rather it is a neutral salt.
Is cabr2 ionic?
If atoms of a molecule are bonded together by electrostatic force, then it is ionic compound means it consists of oppositely charged particles which are tied up with the force. Generally metal and nonmetal form cation and anion respectively to produce this with high melting point.
Calcium is an alkali earth metal categorized in group 2 and Bromine is a non-metal element (halogen) which located in group 17. Following the Octet rule to cover the last filled shell, permanent electron transfer occurs from Ca to Br (not share). Opposite charges packed together.
After forming the ionic bonding the energy of the molecule stabilizes very much, difficult to separate ions as strong Coulombic force is acted between them. Also the charge on cation is (+2), making ions closely packed, results in high melting point of CaBr2 lewis structure around 730 ̊ C.
Monatomic ions are generated specially as electrolytes. “Mono” is a Greek word which here stands for one and “atomic” is the adjective form of atom which is the most tiny part of the matter. Ions are the particles which carry charge over it means number of protons and electrons are not equal.
Sodium (Na) is an element of group 1in the third period of the Periodic table which has one electron in the last filled 3s orbital.
It tends to transfer this electron to get the stability like nearest inert gas Neon with electronic configuration 1s2 2s2 2p6. This makes monatomic cation. As the number of proton in the nucleus are more than the electrons moving in various energetic orbits. This ion called “Sodium plus” or Na+.
During electrolysis from table salt this monatomic cation is also generated in the solution. As it carries charge it can move electricity in the solution.
Aluminum is an element of group number 13 in the third period. 3s and 3p are the last orbitals which have electrons in the atom.
It can remove the three electrons for gaining energetic stability which make a monatomic ion example “Al3+” or “Aluminum 3+”.
K+ as monatomic ion
Potassium is an element in group one of the fourth period. Nearly similar to the nearest inert gas electronic configuration it has one extra electron in 4s orbital.
For energetically stability purpose it remove the electron from the last filled 4s orbital, results in monatomic cation “K+” or “Potassium +”.
Mg2+ as monatomic ion
Magnesium is in 2nd group of the 3rd period with two electrons in the outer filled 3s orbital.
Removing these electrons it makes monatomic cation “Mg2+” which is most important in Chlorophyll formation.
Chlorine belongs to the group 17 of the 3rd period with vacancy of one electron in 3p orbital so that it faces lack of stability.
By accepting one electron it produces monatomic ion example that is “Chloride”. It can be obtain in the electrolysis of table salt.
monatomic ion examples
How monatomic ions form?
Octet rule governs the cause of the formation of ions.
The energy of the orbit of atoms or molecules is most stable when the last occupied orbit fills eight electrons which is also known as Octet rule in atomic structure chemistry. Only group 18 elements in Periodic table have this type of configurational stability.
Except inert gases all atoms in Periodic table have lesser number of electrons in the last filled shell, among which some gain electrons to reach the target or some lose electrons for that. This results in the formation of the single charged atom.
Sometimes molecules especially ionic compounds dissociate and produce simplest elements which have charge over them. Some of the elements belong to the monatomic ion examples. Like NaCl (table salt) dissolves in water and produce two monatomic ions.
Ions can be classified on the nature of the charge it carries. The nature of the charge which the generating ion occupies depends on the electronic arrangement of the outer most filled orbit of that atom.
Formation of monatomic cation
If the last filled orbital of an atom has less number of electrons than its capacity of the half filled shell then the atom tends to transfer the electron to achieve electronic stability. After the transformation the atom has greater number of protons in nucleus than the electrons results cation.
As cations have greater number of protons so they are positively charged elements. Naturally metals form this cations as they have less than or equal to three loosely bounded electrons by nuclear force in the outer orbit. The charge have to be integer number like (+1), (+2) etc.
In a solution these ions with positive charge are attracted towards cathode, these are called cation.
Formation of monatomic anion
If the last filled orbital of an atom has greater number of electrons than the half filled of that orbital can posses, but not full fill the Octet rule then the atom accepts electrons from donor. This acceptance results in formation of anion.
There is greater number of electrons present in the orbit of an anion than the protons in the nucleus of it so the anions are negatively charged particles. The charge of the anion must be integer in number like (-1), (-2) etc.
In solution these are attracted toward the anode, these are called anion.
How to write monatomic ions?
To write this two things have to be considered, one is the atomic symbol of the ion and then have to predict the charge.
The charge gained by the atom after losing or gaining electrons, is shown at the upper right side of the atomic symbol. Sometime in the monatomic ion examples the oxidation state is written in bracket after the atomic symbol of the monatomic ion in Roman numeral.
As an example we can check Lithium as a monatomic ion. This atomic symbol is “Li”. After losing the 2s orbital electron it produce Li+ monatomic ion.
How to name monatomic ions?
There is a particular naming pattern for both positively and negatively charged monatomic ions which are produced from metals or alkali earth metals and non-metals respectively.
For monatomic cation, the name remains same as the name of the element. The positive charge number is attached with the element name. Monatomic ion of Hydrogen is written as “H+” which is called “Hydrogen +”.
To write the monatomic anion a suffix (-ide) has to be added after the element name. Halogen Chlorine forms anion with one atom by taking one electron from donor. To write this monatomic ion (Cl-) we use the name “Chloride”.
Monatomic cation with various oxidation states
For the cationic elements which shows more than one oxidation states means they can be energetically stable in different cationic configuration, are written by their atomic symbol with that oxidation state in Roman numeral in a bracket.
Mainly transitional metals form this type of monatomic ion examples with more than one oxidation states. For the presence of the “d” orbital in these they can form different cationic configuration. For example Iron can be stable in both forms Fe (l) and Fe (ll).
Conclusion
By warping up this article it can be concluded that for gaining energetic stability with full filled orbital, elements form these ions which are monatomic ion examples.
Oxybromide (OBr2) has a central oxygen (O) atom with 6 valence electrons, forming single bonds with two bromine (Br) atoms, each contributing 7 valence electrons. The Lewis structure shows two single O-Br bonds and two lone pairs on the oxygen atom, using 20 valence electrons in total. OBr2 adopts a bent molecular geometry with a bond angle slightly less than 109.5°, influenced by lone pair-bond pair repulsion. The O-Br bonds are polar due to the electronegativity difference (O: 3.44, Br: 2.96), making the molecule polar overall. This structure affects its chemical reactivity and potential use in halogenation reactions.
By the arrangement of the outer most shell electrons, the most energetically stabilized Structure of the obr2 lewis structure can be determined. In a lewis structure loosely bounded electrons by nuclear force are shown around the atomic symbols.
OBr2 Lewis Structure
How to draw lewis structure for OBR2
In obr2 Lewis structure we see a basic diagram where the O-Br bonding skeleton of the molecule is shown.
With using the atomic symbols of Oxygen and Bromine, the obr2 lewis structure shows the position of the atoms in the molecule with excess electrons (beside bonding electrons) around the particular atom.
In modern Periodic table Oxygen is in group16 and Bromine is in group 17. The electrons in the last orbit of Oxygen atom is [He] 2s2 2p4. Electronic arrangement of Bromine atom is [Ar] 3d10 4s2 4p5.
In the outer orbit energy levels of the Oxygen and Bromine atoms, total twenty negatively charged particles (7 + 7 + 6) are available means ten pairs of electron for bond formation in the obr2 lewis structure.
Oxygen is in the center of the obr2 electron dot structure and two Bromine atom connected to the opposite side from the central atom. After this arrangement; valence electrons are placed according to their parent atom.
To cover the last orbit with eight electrons, Oxygen and Bromine atoms share electrons cloud. To make a single bond, two electron cloud are joined, that is shown by a single line in obr2 lewis structure.
In this above obr2 lewis structure, both Oxygen and Bromine atoms have total eight pairs of excess unshared electrons.
OBR2 lewis structure shape
According to the Valence bond theory shape of a molecule is the distorted form of the geometry of the molecule if the central atom contains any unshared electron pair.
The VSEPR theory states with greater charge density the repulsion capability of lone pair electron cloud is greater than the bonding electron pair cloud. The center atom Oxygen has two pairs of non bonding electron cloud which occupies much space to create steric repulsion on the bonding electron cloud.
Oxygen atom in obr2 molecule has sp3 hybrid orbital so the related geometry would be like tetrahedral. But the electron cloud making steric repulsion disturbs the geometry and forms a bent shaped molecule.
OBR2 lewis structure formal charges
Formal charge is the charge density of an atom in a particular electron dot configuration, where it is assumed that the bonding electron cloud is equally distributed.
Formal charge of the atoms in obr2 lewis structure can be calculated by the equation, f = (n – lp – (bp/2)). Where ‘n’ is number of last orbit electrons, lp is unshared electrons and bp is bonding pairs electron of a particular atom.
From the Periodic table, we know that Oxygen atom has six electrons and Bromine atom has seven electrons in their last filled electronic shell. In the obr2 lewis structure each Bromine atom has three lone pairs where Oxygen atom has two pairs of non-bonding electron cloud.
Bromine atom formal charge in above obr2 electron dot structure, {7- 6 – (1/2×2)} = 0.
For Oxygen atom in the lewis structure formal charge is, {6- 4- (1/2×4)} = 0.
As the atoms of this above obr2 electron dot structure have “zero” formal charge, hence this obr2 lewis structure is very much stable in energetically as compared to the other possible structures.
OBR2 lewis structure lone pairs
A pair of electrons which remains unshared in a chemical bond formation is called lone pair. The lone pair electron cloud density only concentrates over containing atom.
To form obr2 lewis structure, twenty last filled orbit electrons of the three atoms are involved. Among these, two pairs (four electrons) of electron result in O-Br bond formation. Eight pairs of electrons are excess, these are lone pair electrons of Dibromine monoxide molecule.
Lone pair electron cloud remains in the control of only one atom, so near to its nucleus than the bonding electron cloud which spread over the average distance of the two bonding atoms. So this negatively charged cloud has high charge density compared to the bonding electron cloud.
That’s why Lone pair cloud is bulky in nature and creates strong steric repulsion. If central atom has the lone electron pair, its geometry becomes distorted.
OBR2 hybridization
According to the Valence bond theory atomic orbitals are intermixes to produce new hybrid orbital with different shape and energy which enhance the capacity of overlapping with other atomic orbital.
In the lowest energy level, Oxygen atom has two paired electrons in 2s and one 2p orbital. It also contains single spin electron in the two 2p orbital. These 2s and 2p orbitals of Oxygen atom intermixes to give sp3 hybridized orbital.
This sp3 hybrid orbital now makes greater overlap to the 4p Bromine orbital.
OBR2 lewis structure resonance
If the electrons which don’t participate in bond formation shift to vacant orbital of adjacent atom then we get a different electron dot structure, which is called resonance structure.
In the obr2 lewis structure both Oxygen and Bromine have extra electrons which not participate in bond formation. For shifting the electron cloud only Bromine atom has vacant 4d orbital in the 4th shell.
From the neutral electron dot structure, one non-bonding electron pair of Oxygen atom delocalizes to the Bromine atom vacant orbital. As the shifting of electron cloud occurs from the central atom to the ligand atom, this is called back bonding.
This rule states that the involving atoms of an electron dot structure orient themselves in a way that they share electron cloud and their outer energy level shell contain eight electrons. This arrangement of the electronic cloud gives stabilization.
Oxygen have six electrons in its outer s and p orbital od 2nd shell, where Bromine atom has seven electrons in the s and p orbital of 4th shell.
They share the electron cloud in such a way that each of them full fills the margin up to eight electrons in the outer shell. Hence the sp3 orbital of Oxygen atom overlap with the 4p orbital of Bromine to satisfy the rule.
OBR2 polar or nonpolar
Dipole moment (µ) causes polarity in a bond, which can be calculated by product of the separated charge (δ) and distance between charges (r) in the bond.
In the Dibromine monoxide molecule, electro-negativity of Oxygen (O) is 3.44 and that of Bromine is 2.96 on the basis of Pauling scale. This difference in the capacity of the bond electron cloud pulling creates charge separation in bond and Dibromine monoxide becomes polar molecule.
If this electro-negativity difference is 0.4 or more than that a charge separation occurs in the bond, which results a dipole, means one end of the bond becomes slightly negative than the other end. This dipole is a vector component directed towards the negative end.
The direction of dipole is towards the atom with greater tendency of pulling the bonding electron cloud, the Oxygen atom. More over this atom has lone electron pairs cloud which also increases the amount of electron dipole towards it.
For the steric electronic repulsion the shape of the molecule becomes bent. Hence the obr2 lewis structure has a net electronic dipole vector which makes the molecule polar.
OBR2 lewis structure bond angle
In an electron dot structure of a molecule the angle between two adjacent bonds (bond angle) can be calculated mathematically by the dot product of vectors.
The central atom Oxygen is sp3 hybridized, so the possible bond angle in the obr2 lewis structures is 109.5̊. But because of the presence of the electron pairs over this atom which don’t take part in bonding, the geometrical angle distorted.
The non-bonding electrons of the Oxygen atom involves in the 2pπ (O) – 4dπ (Br) back bonding with Bromine atom. The lone pair repulsion on the bonding electron cloud lowers. As partial double bond forms, steric repulsion of bonding electron cloud produces.
Bromine is in the 4th period in Periodic table with bulky electronic configuration. In the molecule both Bromine atoms have three pairs of unshared electron cloud. It higher the steric repulsion of bonding electron cloud.
As a whole the bond angle of the molecule increases to 112̊.
OBR2 lewis structure electron geometry
Theoretically we can predict the geometry of obr2 lewis structure by the assumptions of the Valence Shell Electron Pair Repulsion theory.
With the help of the hybridization of central atom, bond length, bond angle, general shape of the molecule we can predict the possible atomic position in a molecule. Various spectroscopic methods like UV-visible, IR, Raman, NMR, Electronic or Mass spectra can help to identify the geometry obr2 lewis structure.
In the obr2 lewis structure sigma (σ) bond is present in the both side of the central Oxygen atom which is sp3 hybridized with two pairs of unshared electron cloud. So this molecule can be generalized as AX2 type molecule with two single bonds and two non-bonding electron pair cloud.
From the VSEPR theory it can be said that this molecule belongs to the tetrahedral geometry group.
OBR2 valence electrons
The last filled energetic shell electrons are called valence electrons which can absorb energy to be in excited state where they can involve in chemical reactions.
In Obr2 lewis structure the three atoms contains twenty that type of electrons. Each bromine atom has seven electrons in the last filled 4s and 4p orbitals whereas the Oxygen atom has that type of six negatively charged particles in the 2nd orbit.
The electrons of the core orbitals in an atomic structure model are tightly bounded to the nucleus by strong nuclear force. With increasing the electronic energy level from the nucleus, the binding force lowers.
For this reason the outer shell electrons are loosely tied up and can move to higher energy level by taking up the provided energy.
OBR2 uses
Dibromine monoxide is used for bromination reactions where the radical form is produced to progress the reaction further.
The compound is only in a stable condition in very low temperature like (-40 ̊) or lower than this. Even it decomposes at room temperature. So not much uses of the compound is known.
The OCS Lewis structure refers to the arrangement of atoms and electrons in the compound known as carbonyl sulfide. This structure is essential in understanding the chemical properties and behavior of OCS. In the Lewis structure, the central atom, carbon, is bonded to both oxygen and sulfur atoms through double bonds. The oxygen and sulfur atoms each have two lone pairs of electrons. This arrangement allows us to visualize the distribution of electrons and predict the molecule’s shape and polarity. Understanding the OCS Lewis structure is crucial in studying its reactivity and its role in various chemical reactions.
Key Takeaways
Atom
Number of Bonds
Number of Lone Pairs
Carbon
2
0
Oxygen
2
2
Sulfur
2
2
Understanding Lewis Structures
Lewis structures are a way to represent the arrangement of atoms and electrons in a molecule. They provide a visual representation of the molecular structure and help us understand the bonding and geometry of molecules. In this article, we will explore how Lewis structures are written, how they show molecular shape, and how to determine if a Lewis structure is polar or nonpolar.
How Lewis Structures are Written
To write a Lewis structure, we follow a set of rules and guidelines. The octet rule is a fundamental concept in Lewis structures, which states that atoms tend to gain, lose, or share electrons in order to achieve a stable electron configuration with eight valence electrons.
Lewis structures are also known as Lewis dot diagrams, as they use dots to represent valence electrons. Each dot represents one electron, and they are placed around the atomic symbol to indicate the electron’s location in the atom.
When writing Lewis structures, we start by determining the total number of valence electrons in the molecule. Valence electrons are the electrons in the outermost energy level of an atom and are involved in chemical bonding. We then distribute these electrons among the atoms, ensuring that each atom has an octet of electrons, except for hydrogen, which only needs two electrons to achieve stability.
In some cases, the octet rule may be violated due to the presence of odd-electron species or atoms with more than eight valence electrons. These exceptions are typically found in molecules with elements from the third period or below.
How Lewis Structures Show Molecular Shape
Lewis structures not only provide information about the arrangement of atoms but also give insights into the molecular geometry. The arrangement of electron pairs around the central atom determines the molecular shape.
The VSEPR (Valence Shell Electron Pair Repulsion) theory is commonly used to predict molecular geometry based on the number of electron pairs around the central atom. According to this theory, electron pairs repel each other, and the molecule adopts a shape that minimizes these repulsions.
The number of electron pairs, both bonding and non-bonding, determines the molecular shape. For example, a molecule with two electron pairs around the central atom will have a linear shape, while a molecule with four electron pairs will have a tetrahedral shape.
How to Determine if a Lewis Structure is Polar or Nonpolar
To determine if a Lewis structure is polar or nonpolar, we need to consider the polarity of the individual bonds and the molecular geometry.
If a molecule has polar bonds, meaning there is an unequal sharing of electrons between atoms, and the molecular geometry is asymmetrical, then the molecule is polar. This means that the molecule will have a positive and a negative end, creating a dipole moment.
On the other hand, if a molecule has nonpolar bonds or symmetrical molecular geometry, then the molecule is nonpolar. In nonpolar molecules, the electron density is evenly distributed, resulting in no dipole moment.
It’s important to note that the presence of polar bonds does not necessarily mean the molecule is polar. The overall molecular geometry must also be taken into account.
Deep Dive into OCS Lewis Structure
OCS, also known as Carbonyl Sulfide, is a chemical compound composed of one carbon atom, one sulfur atom, and one oxygen atom. Understanding the Lewis structure of OCS is crucial in comprehending its molecular geometry, bond angles, and other important properties. In this deep dive, we will explore various aspects of the OCS Lewis structure, including how to draw it, its shape, formal charges, lone pairs, resonance, octet rule, bond angle, and electron geometry.
To draw the Lewis structure of OCS, we need to determine the total number of valence electrons present in the molecule. Carbon has four valence electrons, oxygen has six, and sulfur has six as well. Adding them up, we get a total of 16 valence electrons for OCS.
Next, we arrange the atoms in a way that satisfies the octet rule, which states that atoms tend to gain, lose, or share electrons to achieve a stable electron configuration with eight valence electrons. In the case of OCS, we place the carbon atom in the center, surrounded by the oxygen and sulfur atoms.
To form bonds, we share electron pairs between the atoms. Carbon forms double bonds with both oxygen and sulfur, while oxygen and sulfur each have a lone pair of electrons. This arrangement allows each atom to have a complete octet of electrons.
The Lewis structure of OCS reveals its linear molecular geometry. The linear shape occurs because there are no lone pairs of electrons on the central carbon atom, and the oxygen and sulfur atoms are bonded to it in a straight line. This linear arrangement gives OCS a symmetrical shape.
Formal charges are used to determine the distribution of electrons in a molecule and help identify the most stable Lewis structure. To calculate the formal charges, we compare the number of valence electrons an atom should have (based on its group number in the periodic table) with the number of electrons it actually has in the Lewis structure.
In the Lewis structure of OCS, the carbon atom has a formal charge of zero, while both the oxygen and sulfur atoms have a formal charge of zero as well. This distribution of formal charges indicates that the Lewis structure of OCS is stable.
OCS Lewis Structure Lone Pairs
Lone pairs are pairs of valence electrons that are not involved in bonding. In the Lewis structure of OCS, both the oxygen and sulfur atoms have one lone pair of electrons. These lone pairs contribute to the overall electron density around the atoms and influence the molecule’s properties.
OCS Lewis Structure Resonance
Resonance structures are alternative Lewis structures that can be drawn for a molecule by moving electrons while keeping the positions of the atoms fixed. In the case of OCS, resonance structures can be formed by moving the double bond between the carbon and oxygen atom to the carbon and sulfur atom.
However, despite the presence of resonance structures, the actual OCS molecule does not exhibit resonance. This is because the resonance structures are energetically equivalent and rapidly interconvert, resulting in a hybrid structure that represents the average of the resonance forms.
OCS Lewis Structure Octet Rule
The octet rule states that atoms tend to gain, lose, or share electrons to achieve a stable electron configuration with eight valence electrons. In the Lewis structure of OCS, all atoms satisfy the octet rule. Carbon has a total of eight electrons (four from the double bonds and four from the lone pairs), while both oxygen and sulfur have eight electrons each (six from the double bond and two from the lone pair).
OCS Lewis Structure Bond Angle
The bond angle in the OCS molecule is approximately 180 degrees. This angle arises due to the linear molecular geometry of OCS, where the carbon atom is positioned between the oxygen and sulfur atoms in a straight line. The absence of any lone pairs on the central carbon atom allows for a linear arrangement, resulting in a bond angle of 180 degrees.
OCS Lewis Structure Electron Geometry
The electron geometry of OCS is linear. This geometry describes the arrangement of all electron pairs (both bonding and lone pairs) around the central atom. In the case of OCS, there are no lone pairs on the central carbon atom, and the oxygen and sulfur atoms are bonded to it in a straight line, resulting in a linear electron geometry.
OCS Valence Electrons
OCS, also known as Carbonyl Sulfide, is a molecule composed of one carbon atom, one oxygen atom, and one sulfur atom. To understand its valence electrons, we need to look at the electron configuration of each atom.
Carbon has an atomic number of 6, which means it has 6 electrons. The electron configuration of carbon is 1s2 2s2 2p2. Oxygen has an atomic number of 8, so it has 8 electrons. Its electron configuration is 1s2 2s2 2p4. Lastly, sulfur has an atomic number of 16 and an electron configuration of 1s2 2s2 2p6 3s2 3p4.
To determine the total number of valence electrons in OCS, we need to consider the outermost energy level of each atom. In this case, the valence electrons are the electrons in the 2s and 2p orbitals of carbon, the 2s and 2p orbitals of oxygen, and the 3s and 3p orbitals of sulfur.
Carbon has 4 valence electrons, oxygen has 6 valence electrons, and sulfur has 6 valence electrons. Adding these together, we get a total of 16 valence electrons for OCS.
Now that we know the number of valence electrons, let’s explore the hybridization of OCS.
OCS Hybridization
Hybridization is a concept used to describe the mixing of atomic orbitals to form new hybrid orbitals. In OCS, the carbon atom is bonded to both the oxygen and sulfur atoms. To understand the hybridization, we need to consider the electron pairs and molecular geometry of OCS.
In OCS, the carbon atom forms two sigma bonds, one with the oxygen atom and one with the sulfur atom. Each sigma bond is formed by the overlap of a hybrid orbital from carbon and a hybrid orbital from the respective atom. The oxygen atom also forms a sigma bond with the sulfur atom.
The molecular geometry of OCS is linear, with the carbon atom at the center and the oxygen and sulfur atoms on either side. This linear arrangement is due to the presence of two sigma bonds and no lone pairs of electrons on the central carbon atom.
The hybridization of the carbon atom in OCS is sp hybridization. This means that one s orbital and one p orbital from carbon combine to form two sp hybrid orbitals. These sp hybrid orbitals then overlap with the hybrid orbitals of the oxygen and sulfur atoms to form sigma bonds.
By understanding the valence electrons and hybridization of OCS, we can gain insights into its chemical bonding and molecular structure. These concepts are crucial in understanding the resonance structures, bond angles, and electron density distribution within the molecule.
Properties of OCS Molecule
The OCS molecule, also known as carbonyl sulfide, is a chemical compound composed of one carbon atom, one sulfur atom, and one oxygen atom. It exhibits interesting properties that can be understood by examining its polar or nonpolar nature, Lewis structure molecular geometry, and Lewis dot structure.
OCS Polar or Nonpolar
To determine whether OCS is polar or nonpolar, we need to consider its molecular geometry and the polarity of its bonds. The OCS molecule has a linear molecular geometry, with the carbon atom in the center and the oxygen and sulfur atoms on either side.
The carbon-oxygen double bond and the carbon-sulfur double bond in OCS are both polar bonds due to the difference in electronegativity between the atoms. However, since the molecule is linear and symmetrical, the polarities of these bonds cancel each other out, resulting in a nonpolar molecule overall. Therefore, OCS is considered a nonpolar molecule.
OCS Lewis Structure Molecular Geometry
The Lewis structure of OCS helps us understand its molecular geometry. The Lewis structure of OCS is as follows:
O=C=S
In this structure, the carbon atom is in the center, bonded to both the oxygen and sulfur atoms. The oxygen and sulfur atoms each have a lone pair of electrons. The carbon-oxygen double bond and the carbon-sulfur double bond are represented by lines.
The molecular geometry of OCS is linear, with a bond angle of 180 degrees. This arrangement occurs due to the repulsion between the electron pairs surrounding the central carbon atom. The linear geometry of OCS is a result of the electron pairs and the double bonds.
OCS Lewis Dot Structure
The Lewis dot structure of OCS provides a visual representation of the valence electrons and their distribution in the molecule. In the Lewis dot structure, the valence electrons of each atom are represented by dots.
For OCS, the Lewis dot structure is as follows:
O•=C•=S•
In this structure, each dot represents a valence electron. The carbon atom has four valence electrons, while the oxygen and sulfur atoms each have six valence electrons. The dots are placed around the atomic symbols to indicate the distribution of valence electrons.
The Lewis dot structure of OCS helps us understand the arrangement of electrons and the formation of covalent bonds between the atoms. It also provides insights into the formal charge and electron configuration of the molecule.
By examining the properties of OCS, such as its molecular geometry, electron pairs, and covalent bonding, we can gain a deeper understanding of its chemical structure and behavior. The combination of the octet rule, Lewis dot diagram, VSEPR theory, and molecular orbital theory allows us to analyze the electron density and atomic orbitals involved in the formation of OCS. Overall, OCS is a nonpolar molecule with a linear structure, making it an intriguing compound to study.
Practical Applications of OCS
OCS Uses
OCS, or Carbonyl Sulfide, has several practical applications in various fields. Let’s explore some of its uses:
Industrial Applications: OCS is commonly used as a precursor in the production of thiocarbamate pesticides, which are widely used in agriculture to control pests and protect crops. It is also utilized in the manufacturing of rubber chemicals and pharmaceuticals.
Environmental Monitoring: OCS plays a crucial role in studying the carbon cycle and atmospheric chemistry. Its presence in the atmosphere can be measured to gain insights into the exchange of carbon dioxide between the land, oceans, and the atmosphere. This information helps scientists better understand climate change and global warming.
Laboratory Research: OCS is frequently used as a reagent in chemical laboratories for various experiments. Its unique properties make it a valuable tool in the synthesis of organic compounds and the study of chemical reactions.
Medical Applications: OCS has shown potential in medical research as a possible treatment for certain diseases. Studies have indicated that it may have antimicrobial properties, making it a candidate for developing new antibiotics. Additionally, OCS has been investigated for its potential use in cancer treatment due to its ability to inhibit the growth of tumor cells.
OCS Lewis Structure in Real-World Context
To understand the practical applications of OCS, it is essential to examine its Lewis structure and how it relates to real-world scenarios.
In the Lewis dot diagram of OCS, the central atom is carbon (C), surrounded by an oxygen atom (O) on one side and a sulfur atom (S) on the other. The carbon atom has a double bond with the oxygen atom and a single bond with the sulfur atom. This arrangement follows the octet rule, where each atom aims to have a full outer shell of electrons.
The molecular geometry of OCS is linear, with bond angles of approximately 180 degrees. This structure is determined by the VSEPR theory (Valence Shell Electron Pair Repulsion), which predicts the arrangement of electron pairs around the central atom.
The Lewis structure of OCS is also important in understanding its chemical bonding and molecular structure. It helps us identify resonance structures, lone pairs, and the overall electron density distribution within the molecule. These factors influence the molecule’s reactivity and physical properties.
In terms of hybridization, the carbon atom in OCS undergoes sp hybridization, forming sigma (σ) bonds with both the oxygen and sulfur atoms. The double bond between carbon and oxygen consists of one sigma bond and one pi (π) bond, while the single bond between carbon and sulfur is a sigma bond.
The understanding of OCS’s Lewis structure and molecular geometry is crucial in various fields. For example, in organic chemistry, it helps chemists predict the behavior of OCS in reactions and design new compounds. In atmospheric science, it aids in the analysis of OCS concentrations and its impact on climate change.
Comparison with Other Lewis Structures
OCS vs H2O2 Lewis Structure
When comparing the Lewis structures of OCS and H2O2, we can observe some interesting differences. OCS, also known as carbonyl sulfide, consists of one carbon atom bonded to one sulfur atom and one oxygen atom. On the other hand, H2O2, or hydrogen peroxide, contains two hydrogen atoms bonded to two oxygen atoms.
In terms of the octet rule, both OCS and H2O2 satisfy this rule. Each atom in both molecules has a complete octet of valence electrons, which allows them to be stable. However, the molecular geometries of these two compounds differ due to the arrangement of their electron pairs.
In OCS, the central carbon atom is surrounded by two electron pairs and two bonding pairs. This gives OCS a linear molecular geometry. On the other hand, H2O2 has a bent molecular geometry due to the presence of two lone pairs on each oxygen atom. These lone pairs repel the bonding pairs, causing the molecule to adopt a bent shape.
OCS vs SOCl2 Lewis Structure
Now let’s compare the Lewis structures of OCS and SOCl2. OCS, as mentioned earlier, consists of one carbon atom bonded to one sulfur atom and one oxygen atom. SOCl2, or thionyl chloride, contains one sulfur atom bonded to two chlorine atoms and one oxygen atom.
Both OCS and SOCl2 satisfy the octet rule, as all atoms have a complete octet of valence electrons. However, the molecular structures of these compounds differ due to the arrangement of their electron pairs and the presence of multiple bonds.
In OCS, the central carbon atom is bonded to the sulfur atom through a double bond, while the oxygen atom is bonded to the carbon atom through a single bond. On the other hand, in SOCl2, the central sulfur atom is bonded to the oxygen atom through a double bond and to the two chlorine atoms through single bonds.
OCS vs OCN- Lewis Structure
Lastly, let’s compare the Lewis structures of OCS and OCN-. OCS, as we have discussed, consists of one carbon atom bonded to one sulfur atom and one oxygen atom. OCN-, or cyanate ion, contains one carbon atom bonded to one nitrogen atom and one oxygen atom, with an extra electron giving it a negative charge.
Both OCS and OCN- satisfy the octet rule, as all atoms have a complete octet of valence electrons. However, the presence of the negative charge in OCN- affects its molecular structure. The extra electron in OCN- creates a lone pair on the nitrogen atom, which influences the molecular geometry.
In OCS, the central carbon atom is bonded to the sulfur atom through a double bond, while the oxygen atom is bonded to the carbon atom through a single bond. On the other hand, in OCN-, the central carbon atom is bonded to the nitrogen atom through a triple bond, and the oxygen atom is bonded to the carbon atom through a single bond. The presence of the lone pair on the nitrogen atom causes OCN- to have a linear molecular geometry.
Frequently Asked Questions
What is the Lewis structure of H2O2 and its electron geometry?
The Lewis structure of H2O2 (Hydrogen Peroxide) involves the formation of single bonds between each Hydrogen (H) and Oxygen (O) atom, and a single bond between the two Oxygen atoms. The remaining electrons on the Oxygen atoms form lone pairs. The electron geometry of H2O2 is bent or V-shaped, due to the presence of two bonding pairs and two lone pairs on each Oxygen atom.
How does the Lewis structure show the molecular shape of a compound?
The Lewis structure provides a visual representation of the arrangement of atoms in a molecule and the bonding between them. It shows the location of valence electrons, which participate in bond formation. The shape of the molecule can be predicted using the Valence Shell Electron Pair Repulsion (VSEPR) theory, which states that electron pairs around a central atom orient themselves as far apart as possible to minimize repulsion.
How can I determine if a Lewis structure is polar or nonpolar?
Polarity in a molecule depends on the difference in electronegativity between the atoms and the symmetry of the molecule. If there is a significant electronegativity difference, the bond is polar. However, if the molecule is symmetrical, the polarities of the bonds may cancel out, resulting in a nonpolar molecule. For example, in the OCS Lewis structure, although the bonds are polar, the linear shape of the molecule makes it nonpolar.
How are Lewis structures written?
Lewis structures are written by first determining the total number of valence electrons in the molecule. Then, the atoms are arranged with the least electronegative atom usually in the center. Bonds are formed by sharing electron pairs between atoms, and the remaining electrons are distributed as lone pairs. The structure should satisfy the Octet Rule, where each atom (except Hydrogen) is surrounded by eight electrons.
What is the resonance in the OCN- Lewis structure?
In the Lewis structure of the OCN- ion, resonance occurs because there are multiple valid ways to arrange the double and single bonds between the Oxygen, Carbon, and Nitrogen atoms. These different structures, called resonance structures, contribute to the overall structure of the ion, making it more stable.
What is the formal charge in the OCN- Lewis structure?
The formal charge is calculated by subtracting the number of valence electrons in the isolated atom from the number of valence electrons assigned to it in the molecule. In the OCN- Lewis structure, the formal charges are -1 for Oxygen, 0 for Carbon, and 0 for Nitrogen.
What is the molecular geometry of the OCS Lewis structure?
The molecular geometry of the OCS (Carbonyl Sulfide) molecule is linear. This is determined using the VSEPR theory, which takes into account the number of bonding pairs and lone pairs around the central atom.
What is the Lewis structure of SOCl2 and its molecular geometry?
The Lewis structure of SOCl2 (Thionyl Chloride) involves a double bond between the Sulfur (S) and Oxygen (O) atoms, and single bonds between the Sulfur and the two Chlorine (Cl) atoms. The remaining electrons on the Sulfur atom form a lone pair. The molecular geometry of SOCl2 is a seesaw shape, due to the presence of three bonding pairs and one lone pair on the Sulfur atom.
How does the Lewis structure of CH3COO- show resonance?
The Lewis structure of the CH3COO- ion (Acetate) shows resonance because there are two equivalent structures that can be drawn with the double bond between the Carbon and one of the Oxygen atoms. These two structures are resonance forms and contribute equally to the overall structure of the ion.
What is the shape of the HCN Lewis structure?
The shape of the HCN (Hydrogen Cyanide) molecule is linear. This is determined using the VSEPR theory, which takes into account the number of bonding pairs and lone pairs around the central atom.
In a human body, twenty types of Amino acids are found. This article describes the Valine amino acid structure and its properties.
Some α-amino acids which human body can’t bio-synthesize, are known as essential amino acids. valine is an essential amino acid. The chemical name of Valine is 2-amino-3-methylbutanoic acid. Valine amino acid structure is non-polar in nature.
Structure of Amino acid: Valine
Amino acids are carboxylic acids with amino group. Molecular formula of valine is C5H11NO2. It is denoted as “Val”.
Amino acids have commonly two main functional groups. Among these one group is amino group (-NH2) and another functional group is Carboxyl group (-COOH). Beside these the amino acid structures have a side chain attached to the mid-point Carbon atom. The chain is shown as ‘R’ for general structure.
(-NH2) group and the (-COOH) group are common in valine amino acid structure like other amino acids. A hydrocarbon group (isopropyl group) is the “R” group in the Valine structure. There also an α-Hydrogen atom is attached to the Carbon atom of Valine.
Valine as an alpha (α) amino acid
The alpha (α) amino acids are basic component of protein molecules.
The center Carbon atom of Amino acid is connected with both the amino group and the carboxyl group, the Amino acid is named as alpha (α) – Amino acid.
In Valine structure both (-NH2) and (-COOH) groups are attached to the one Carbon atom in the middle of the molecule. Hence it can be referred to the α-amino acid.
Amino acid structure
Zwitterion of Valine amino acid structure
If a molecule carries both opposite charged groups that is positive and negative in equal number, this molecule is identified as a zwitter ion.
Amino acids has both (-NH2) and (-COOH) groups. (-COOH) group is acidic in nature so it can transfer the Hydrogen atom. This H+ ion is accepted by the Nitrogen atom of amino group as it is basic in nature. Hence Valine exists as a zwitter ion with both opposite charged groups.
As a dipolar ion, positive and negative charges are present in the Valine molecule. Both the contrasting charges cancel out each other. As a whole the Valine molecule remains as neutral molecule.
zwitterion of valine amino acid structure
The amino acids have charge separation present in the zwitter ion form. So Valine can conduct electricity for the configuration with charges in the solid phase.
pH dependence of Valine structure
The stability of the dipolar ion of the amino acid depends on the pH of the solution.
At a lower pH (acidic pH), the ion accepts H+ ion from solution and forms positive charged ion. In a higher pH (strong basic pH), the ion transfers hydrogen ion (H+) and forms a molecule with negative charge.
At a particular pH value an amino acid remains as a zwitterion which is called the isoelectric point of that amino acid in that solution. For Valine the isoelectric point is 6.0.
If a molecule or a part of it resists contact with the water molecule, then the nature is called hydrophobic nature.
Valine has isopropyl group in the branched hydrocarbon side chain of the center atom. It is a non-polar group. This makes it hydrophobic in nature.
Solubility of Valine amino acid
Solubility of Valine depends on the side isopropyl chain, pH of solution etc.
As Valine structure carries charged groups like carboxylaate and aminium, it can soluble in water at a certain temperature.
But Valine can’t soluble in non-polar solvents, like Benzene. Because these solvents which are not polar can’t able to dissolve the ionic part of Valine molecule.
Chiral Carbon of Valine molecule
Carbon atom can make four bonds in a molecule. If the four groups or atoms are distinct; then the Carbon atom is named as a chiral Carbon center.
A molecule with chiral carbon has optical rotatory power. In order to arrangement of the groups in the molecule, it classifies as ‘D’ and ‘L’ isomer. These structures are carbon copy with respect to each other with same molecular formula but not super impose.
Valine has four distinct groups around the Carbon in the center. So Valine is optically active molecule. It has two individual structures like D- Valine and L- Valine.
D- valine amino acid sttructure
The planer structure of Valine with D-configuration has the amino group (-NH2) to the right side of the Carbon atom. The Valine with L-configuration has the amino group to the left side of the Carbon atom. The Hydrogen atom (α) is in the horizontal position in both structures.
L- valine amino acid structure
Peptide bond formation
Amino acids are the basic component in protein structure.
Amino acids are joining one after another like blocks to produce protein molecules. The new amide bond which formed is called the Peptide bond.
The amide bond is occurs between the carboxyl group and amino group of two adjacent Amino acid. The left side of long chain is called N- terminus which have a free amino group and right side of the chain is called C- terminus with a free carboxyl group.
Valine also participates in peptide bond formation to produce proteins.
Dichlorine monoxide (OCl2) has a central oxygen (O) atom with 6 valence electrons, bonded to two chlorine (Cl) atoms, each contributing 7 valence electrons. The Lewis structure shows two single O-Cl bonds and two lone pairs on oxygen, using 20 bonding electrons. The molecule adopts a bent geometry with a bond angle slightly less than 109.5°, due to lone pair-bond pair repulsion, typical of sp³ hybridization. The O-Cl bonds are polar due to the electronegativity difference (O: 3.44, Cl: 3.16), making the molecule polar overall. This structure affects its reactivity and properties, like being a reactive oxidizing agent.
Outer shell electrons of the involved atoms are shown in the ocl2 lewis structure. These electrons effect the properties of the molecule. So this ocl2 lewis structure gives explanation about the hybridization, geometry, shape and other things of the molecule.
How to draw lewis structure for OCL2
OCl2 Lewis structure gives idea about how the molecular geometry should be.
In the ocl2 lewis structure, the atomic symbols of Oxygen atom is written as ‘O’ and for Chlorine atom it is written as ‘Cl’. The outer shell electrons of atoms are represents by dot symbols around the atoms in the dichlorine monoxide molecule.
Oxygen is in group number 16 and Chlorine is in group number 17 in the modern Periodic table. Oxygen atom electronic configuration is 1s2 2s2 2p4. Chlorine atom electronic configuration is [Ne] 3s2 3p5. So Oxygen atom has six negative charged particles and Chlorine atom has seven negative charged particles in the last electronic shell.
Hence the total number of the negative charged particles in the last orbit is {6 + (7×2)} = 20 in the molecule. So ten pair of electron take part in molecule formation.
Oxygen atom occupy center of the molecule. Two Chlorine atoms are in the both side of the central atom. Two electron pairs form sigma bonds (single bond) between the Oxygen atom and the two Chlorine atoms. Remaining eight electron pairs are unshared electrons.
The sigma (σ) bonding electron cloud between the Oxygen and Chlorine atoms in the OCl2 molecule is drawn by single straight line. Central atom Oxygen has four unshared electrons that is two pairs of electron cloud.
OCL2 lewis structure shape
VSEPR theory (Valence shell electron pair repulsion theory) can help to understand the facts about shape of ocl2 lewis structure.
Minimum repulsion between electron cloud of the last orbit, only can stabilize the structure of a molecule. So the outer shell electron cloud of OCl2 molecule spread over the molecule accordingly.
This theory assumes that, the electrons which don’t participate in bonding formation (lone pair) remain over the originate atom. These spread over more space than the electron cloud of bonding.
For this the repulsion between unshared electron cloud is greater than the electron cloud of bonding pair. Geometry of the molecular structure is disarranged by this.
After atomic orbital mixing up of the central atom Oxygen, new sp3 hybridized orbital occurs in the molecule. This hybridization of atomic orbital refers the tetrahedral molecular geometry. But central atom Oxygen atom has two pairs of unshared electrons.
For the repulsion, the structure of the OCl2 molecule becomes disturbed. To settle the steric repulsion, the angle between the two bonds decreases. Hence shape of the molecule changes to “V” liked shape.
If a lewis dot structure has “0” formal charge, it is the structure which largely contribute in the molecule configuration.
Formal charge = (total number of last shell electrons –unshared electrons – (1/2× electrons in bonding pairs))
From the Periodic table, we know that number of last occupying shell electrons of Oxygen are six and Chlorine has seven electrons. Oxygen atom has four electrons which don’t involve in sigma (σ) bonding whereas each Chlorine atom has six that type of electrons.
Oxygen atom formal charge = {6- 4- (1/2×4)} = 0.
For each Chlorine atom formal charge = {7- 6- (1/2×2)} = 0.
As the atoms of the molecule have “0” (zero) formal charge, hence the formal charge of dichlorine monoxide (OCl2) is also Zero.
OCL2 lewis structure lone pairs
Lone pair contains the electrons which don’t involve in the bonding. These electron pair cloud remains over the parent nucleus.
In OCl2 lewis structure, there twenty outer shell electrons are present. Among these four electrons (two pairs) participate in bonding pair of O-Cl. Remaining all sixteen electrons are non-bonding electronswhich are called lone pair.
From the orbital hybridization, we know that lone pair electron cloud of Oxygen atom belongs to the 2s and 2p orbital. Besides this each Chlorine atom has 3 pairs of electrons which don’t participate in bonding in the OCl2 molecule.
The shape of the molecule is dependent on these non- bonding electrons or lone pairs.
OCL2 hybridization
Hybridization is an important concept as it describes the geometry and shape of a molecule.
Atomic orbitals are fused together in hybridization.The new hybridized orbital has different energy, shape than it has before. In the dichlorine monoxide molecule, we consider the hybridization of the Oxygen atom. Hybridization will help to understand the molecular structure.
In the ground state energy level, Oxygen atom has two 2p orbital electrons which don’t pair up. So these electrons can form two sigma bonds with Chlorine atoms. Oxygen atom has four 2s and 2p orbital electrons which pair up. These electrons can’t participate in bonding.
After hybridization of the outer orbital, orbital overlapping becomes greater in bond formation. The new hybridized Oxygen orbital of OCl2 molecule is sp3 in nature. This sp3 Oxygen orbital is mixing up with the Chlorine 3p orbital to form molecule.
OCL2 lewis structure resonance
OCl2 lewis structure may have many resonance forms. The resonating electron dot structures form by delocalizing of the nonbonding electrons.
In dichlorine monoxide molecule, both the Oxygen and Chlorine atoms have pair of nonbonding electrons. The electron cloud can shift to produce resonating structure. Among these structures, which formal charge will be ‘zero’ that would be the greater contributing configuration.
Chlorine atom has vacant 3d orbital in the outer shell. So it can accept the negative charge from Oxygen atom easily. As central atom, Oxygen accepts electron so Pi (π) back bonding occurs i.e.2pπ (O) – 3dπ (Cl) in the molecule.
It means that the electron cloud transfer from Oxygen atom to the available vacant Chlorine atom 3d orbital. Hence partial pi (π) bonding generates in the molecule and we get the resonating structure of dichlorine monoxide molecule.
The resonating structure of the OCl2 molecule has a great effect on the bond angle and shape of the molecule.
The electronic stability of the orbital of a molecule can be described by this Octet rule. This rule states the maximum capability of electrons in outer orbit of atom.
All main group elements of the Periodic table favour eight electrons in the last energetic shell. So name of the rule is Octet rule. Only Hydrogen atom can’t maintain the rule. Its orbit contains only “s” orbital with maximum capability of two electrons.
Oxygen atom has six electrons in the 2s and 2p orbital. It tends to accept two more electrons. Chlorine atom has seven electrons in the 3s and 3p orbital, so it accepts one electron from donor atom to become stable like Argon gas.
In this way both Oxygen and Chlorine cover the outer most orbits with eight electrons. For this full filled orbit, the energy of the molecule lowers. The configuration of the molecule stabilizes.
OCL2 polar or nonpolar
Polarity is observed in a molecule if the molecule has polar bonds. Bond dipole moment (µ) can be calculated by multiplying the separated charge (δ) and distance between charges (r).
In the dichlorine monoxide molecule, electro-negativity of Oxygen (O) is 3.44 and that of Chlorine is 3.16 on the Pauling scale. An unequal charge separation occurs in the molecule for this difference between the electro-negativity of the atoms. For this a net dipole moment generates and OCl2 becomes polar molecule.
The electro-negativity of an atom is measured by the potential that how strongly it can pull the electron cloud of a covalent bond towards it. It is a direction vector component.
From the hybridization we know the Oxygen atom has two lone pair electron cloud with greater electro-negativity. It has a particular direction vector of dipole moment. Two Chlorine atoms have also specific direction of dipole moment.
From VSEPR theory we know that for steric repulsion, the shape becomes “V” liked. For the bent shape, a net direction vector of dipole moment remains. That’s why OCl2 is polar molecule.
The angle between the bonding electron cloud is called bond angle which plays important role in the shape of the molecule.
As the molecule is sp3 hybridized, the theoretically calculated bond angle should be 109.5 ̊. But for the repulsions of electron cloud present in the molecule the bond angle becomes 110.9 ̊.
The unshared electrons of the Oxygen atom form the 2pπ (O) – 3dπ (Cl) back bonding. The steric repulsion between these electron cloud decreases. As a result of the formation of the partial pi bond (double bond), the repulsion between the bond pairs electron cloud increases.
As a result we get greater bond angle than the calculated one.
OCL2 lewis structure electron geometry
The VSEPR theory (Valence shell electron pair repulsion theory) can direct to find the geometry of OCl2 lewis structure.
From the hybridization, we know that the molecule has two bond pair and the central atom has two non-bonding electron pair cloud. So ocl2 molecule has tetrahedral geometry.
Geometry of a lewis structure is an ideal structure considered for the molecule. But actually it does not exist for facts like repulsion of the electron cloud in the molecule. The disturbed structure is called shape of the molecule.
Geometry is the arrangements of the bonding electron pairs in the three dimensional space. OCl2 is sp3 hybridized. For the four main electron pairs the ideal geometry will be tetrahedral.
OCL2 valence electrons
The electrons which occupy the outer orbit energetic state of an atom are called the valence electrons of that atom.
OCl2 lewis structure has twenty outer shell electrons. These electrons participate in the molecule formation and also responsible for the chemical properties of the molecule.
Oxygen has six electrons in the last energy shell and Chlorine atom has that of seven. So one Oxygen and two Chlorine atoms have twenty electrons in the outer orbits of the atoms. Hence the number of the valence electrons in the molecule is twenty.
OCL2 uses
OCl2 is an inorganic compound. It is soluble in water and as well as in the organic solvents.
The molecule has both Oxygen and Chlorine atom. For the presence of the Oxygen it is an oxidizer molecule. For the Chlorine atom it is a chlorinating agent also. It can be used in the chlorination process for the water purifying for drinking. It can be used in the ring or side- chain chlorination of any organic reaction.
