Class 11 Chemistry includes various important concepts like chemical bonding that will be helpful throughout your higher education if you plan to be a part of the STEM field. It is also useful in the preparation of NEET, JEE, and other entrance exams in India, especially, in Class 11 the chapter on Chemical Bonding and Molecular Structure. The lengthy description of the concepts and theories makes it look much more difficult than it is. Read these to-the-point study notes on Class 11 Chemical Bonding and Molecular Structure and make this chapter your strength!
This Blog Includes:
- Chemical Bond
- Kossel Lewis Approach to Chemical Bonding
- Lewis Symbol
- Ionic Bond
- Favorable factors for forming ionic bonds
- General Characteristics of Ionic Compound
- Method of Writing Formula of Ionic Compound
- Bonding Between Atoms
- Lewis dot structure
- Limitation of Lewis dot structure
- Born Haber Cycle
- VSEPR model
- The Valence bond approach
- Covalent Bond
- Polar Covalent Bond
- Properties of Covalent Compound
- Octet Rule
- Polyatomic molecules
- Molecular Orbital Theory
- Electronic Configuration
- Bond Length
- Bond Angle
- Bond Order
- Bond Enthalpy
- Dipole Moment (μ)
- Hydrogen bond
- Metallic Bond
- Fajan’s Rule
- Exercises – Chemical Bonding and Molecular Structure class 11
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Class 11 chemistry NCERT Solutions explains chemical bonds as the attractive forces that hold the various chemical constituents together in different chemical species. Bonds are formed to get stability with a surplus release of energy.
Credits – Britannica
Kossel Lewis Approach to Chemical Bonding
It states that atoms take part in bond formation for completing the octet or for acquiring the electronic configuration of its nearest inert gas atoms. This is achieved by gaining, losing, or sharing electrons.
It is a method of denoting a valence electron where the chemical symbol of the element is surrounded by dots.
Also read: Scope in Chemistry
It is the chemical bond formed due to the complete transference of electrons from one atom to another. Such that each atom attains the configuration of its nearest noble gas. The solutions provided explain that there are three favorable factors for the formation of an ionic bond:
- Metals must possess the lowest ionization enthalpy
- Non-metals must have the highest electron gain enthalpy
- The energy released during lattice enthalpy must be high
Other concepts which are explained are the three-dimensional aggregation of negative and positive ions in an ionic bond, also known as the crystal lattice. It further explains the charge balance in a crystalline solid between negative and positive ions popularly known as the enthalpy of lattice formation.
Image Credits: Hyperphysics
Favorable factors for forming ionic bonds
- The ionization enthalpy of metals have to be at their lowest
- Electron gain enthalpy of nonmetals should be at their highest
- Lattice enthalpy should also be high
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Atoms which are carrying either negative or positive charges are termed ions. Atoms that have positive charges are called cations whereas those that are carrying negative charges are termed anions. Anions are usually formed by non-metals whereas cations are usually formed by metals.
General Characteristics of Ionic Compound
- Ionic compounds are usually sold in nature
- They have high boiling and melting points
- Ionic compounds can get dissolved in polar solvents like water but they are insoluble in non-polar solvents such as C014, benzene, etc.
- Ionic compounds exhibit good conductivity in an aqueous and molten state
- They have a structure that resembles a crystal
Method of Writing Formula of Ionic Compound
- Write the cation symbol on the left and the anion sign on the right.
- On the top of each symbol, write their electrical valencies in figures like AXBY.
- HCF is used to divide their valencies.
- Apply the criss-cross rule to the formula, which is AyBx.
- Al2(So4)3 is the formula for aluminum sulfate.
Do Refer: Class 11 Thermodynamics
Bonding Between Atoms
The solutions further explain in detail the formation of a covalent bond. A covalent is formed when two atoms share an electron pair between themselves whereas multiple bonds are formed as sharing of more than two electron pairs. Some bonded atoms have an additional pair of electrons, which are not involved in bonding known as lone pairs of electrons.
Lewis dot structure
Our solutions further explain the Lewis dot structure. It shows the arrangement between lone pairs and bonded pairs around each atom in a molecule. Further, it discusses the crucial parameters associated with chemical Bonding like bond length, bond angle, bond order, bond enthalpy, and bond polarity.
Limitation of Lewis dot structure
The solution goes on to explain that a single Lewis structure cannot describe the number of molecules and polyatomic ions. The number of descriptions and skeletal structure written taken together represents ions or molecules. It is a fundamental concept of chemical bonding known as the concept of resonance. A resonance hybrid, which is the canonical form taken together, represents the ion or molecule.
Let Us Have a Look at the Study Notes Of Class 11 Redox Reactions!
Born Haber Cycle
The cycle has its basis on the fact that the formation of the ionic compound could be the outcome of a direct combination of the elements by an alternate process in which:
- The reactants are vaporized for getting converted into a gaseous state
- The gaseous atoms are converted into ions
- Ionic lattice is formed through the combination of gaseous ions
It is used for predicting geometrical shapes between molecules. Also postulates that pairs of electrons repel each other and therefore tend to remain as far as possible. It states that molecular geometry is determined through the repulsions between lone pairs and bonding pairs. The order of repulsion can be portrayed as lp-lp->lp-bp->bp-bp.
The Valence bond approach
Our solutions further explain that a covalent bond is concerned with the energetics of covalent bond formation, which is leftover in the Lewis and Vesper models. Valence Bond approach discusses bond formation in terms of orbital overlaps. For instance, the formation of two-hydrogen molecule H2 from 1s orbital overlaps when they are singly occupied. As the two come close to one another, there is a decrease in the energy of the system. The energy of the system hits its lowest at the equilibrium of internuclear distance. Any further attempt to bring the nuclei closer causes a sudden increase in energy and destabilizes the molecule. It is due to orbital overlap that the density between the nuclei increases and helps in bringing them closer. However, overlap alone does not typify bond length and bond enthalpy; there are other factors too.
A covalent bond is a chemical link created between two atoms by mutual sharing of electrons to complete their octets or duplets. The quantity of electrons supplied by each atom is known as covalency. For example, the creation of CI2/
Polar Covalent Bond
When two atoms form a covalent connection, the shared pair is shifted towards the electronegative atoms, resulting in a denser electron concentration surrounding the more electronegative atom. Ionic characteristics can arise in such covalent connections, which are also known as polar covalent bonds. A typical example is the H-Cl.
Properties of Covalent Compound
- At room temperature covalent compounds exist in gaseous and liquid states as a result of the magnitude of the intermolecular force.
- The boiling and melting points of covalent compounds are low
- They have low conductivity of electricity due to the absence of free ions or electrons for electrical conductivity
- They are soluble in non-polar solvents such as benzene but insoluble in water.
As per the Octet rule when a covalent bond is formed the atoms attain the electronic configuration of inert gas. This state is reached by the atom either by losing or sharing electrons with other atoms.
To explain the characteristic of polyatomic molecules, Pauling introduced a concept known as the hybridization of an atomic orbital. These are represented as sp3, sp2, and sp of atomic orbital of B, C, Be, N, O. These are put to use for explaining the geometrical shape and form of molecules like BeCl2, NH3, BCI3, CH3, and H2O. It also goes on to explain the formation of multiple bonds, such as C2H2 and C2H4.
Molecular Orbital Theory
The solutions provided describe in detail the molecular orbital theory that highlights the bonding in terms of the arrangement and combination of atomic orbital associated with a molecule as a whole. The number of the molecular and atomic orbital is always equal, from which they are formed. The bonding of molecular orbital surges the density of electrons among the nuclei and has lower energy in comparison to individual atomic orbital. Anti-bonding molecular orbital possesses a region of zero electron density among the nuclei and possesses more energy in comparison to the individual atomic orbital.
The solution goes on to explain the representation of electronic configuration through the filling of electrons in the molecular orbital. Here the Hund’s rule and Pauli’s exclusion principle are put into use. The number of electrons in bonding molecules determines the stability of molecules, which is greater than the anti-bonding molecular orbital.
The distance between the two nuclei of a covalently bonded molecule is referred to as bond length. There is an increase in bond length when the size of the bonded atoms increases. Likewise, it decreases with an increase in the number of bonds between bonded atoms.
The length of a bond can be determined via X-ray diffraction or the electron differences approach.
Also Read: Class 11 Hydrocarbons
Molecules that are covalently bonded have more than two atoms and it forms an angle with each other. Technically, this angle is called the bond angle. When the size of the central atoms increases the bond angle decreases. The two common factors that cast an effect on the bond angle are Hybridization of the central atom and lone pair repulsion. X rays diffraction method is used to determine it.
It is termed the number of covalent bonds present in a molecule. NUmerically it could be explained as
Bond Order= ½[ electrons present in the bonding orbital – electrons present in the antibonding orbitals ]
When one mole of covalent bond is formed a certain amount of energy is released which is termed as bond dissociation enthalpy. You can consider it as the amount of energy necessary for breaking one mole of bonds of the same nature for separating bond atoms in the gaseous state.
It has to be remembered that bond dissociation enthalpy and bond enthalpy are equal in magnitude but opposite in sign. With the increase in bond order, there is also an increase in bond enthalpy and a decrease in bond length. Factors that affect bond enthalpy are:
- atomic size
- extent of overlapping
- bond order
Dipole Moment (μ)
It is defined as the product of the charge magnitude and the distance between the positive and negative charge centers.
μ = charge (Q) x distance of separation (r)
Dipole Moment is expressed in Debye. (D).
1 D = 1 * 10-18 esu-cm = 3.33564 * 10-30 C-m
where c is coulomb and m is meter.
(The shift in electron density is symbolized by broken arrow)
NH3 has a higher dipole moment than NF3.
Resultant dipole moment.
μ = √μ21 + μ22 + 2μ1μ2 cos θ
The formation of a hydrogen bond takes place when a hydrogen atom is placed between two highly electronegative atoms like F, O, and N. They can be intermolecularly known for their existence between two or more molecules of the different or same substances or intramolecular, i.e. present between the same molecule.
Hydrogen bonds tend to have a powerful effect on the properties and structure of many compounds. You can define a Hydrogen bond as the attractive force binding hydrogen atoms of one molecule with an electronegative atom of another molecule such as F, O, or N.
H bonds are of two types: Intramolecular hydrogen bond and Intermolecular Hydrogen bond. The physical state of the compound determines the magnitude of H bonding. It is minimum in the gaseous state and maximum in a solid-state. An intermolecular Hydrogen bond is formed between molecules of the same or different compound. For example, the HF molecule. Whereas, an Intramolecular Hydrogen bond is formed by two highly electronegative atoms present in the same molecule.
The force of attraction between a metal ion and a group of electrons inside its sphere of influence is known as the metallic bond. The electron-sea hypothesis of metallic bonding explains a variety of metal characteristics.
Strength of bonds
Ionic bond > covalent bond > metallic bond > H-bond
Ionic bonds do have a partial covalent character which was discussed by Fajan. Consider a bond between smaller cations and the size of the anion is larger in such a situation an ionic bond expresses stronger covalent characteristics. If the cation or anion has a higher charge in such a situation the ionic bond exhibits a greater covalent character.
Also Read: Class 11 Hydrocarbons
It states that a single Lewis structure cannot explain all the properties of molecules. The molecule may exist in different structures each of which is typical of the underlying properties. It is said that the actual structure lies somewhere between all structures contributing are also called resonance hybrids and the various individual structures are termed as canonical structures or resonating structures. Resonance is the common term used for describing this phenomenon.
Since the energy of the resonance hybrid is lesser in comparison to the single canonical structure so it is considered that resonance stabilizes the molecule. The term resonance energy is used to denote the energy gap between the resonance hybrid and its most stable contributing structure. As the resonance energy of the molecule increases so does its stability.
Molecules that show resonance bond order can be mathematically expressed as quotients of the total number of bonds existing between two atoms in all the structures divided by the total number of resonating structures.
Enjoyed Reading Chemical Bonding And Molecular Structure Class 11 Also Refer NCERT Chemistry Class 11
Exercises – Chemical Bonding and Molecular Structure class 11
- Describe the formation of a chemical bonding.
- Write down the Lewis dot structure for the following elements: Br, Mg, N, B, and O.
- Pen down the Lewis symbol for the following ions and atoms: H and H, Al and Al3+, S, and S2.
- Draw the Lewis structure for the following molecules and ions: HCOOH, SiCl2, H2S, CO23, BeF2.
- Write down the favorable factors for the formation of an ionic bond.
- Using the VSEPR model, please discuss the shape of BeCl2, BCl3, AsF3, PH3, SiCl2.
- Discuss the method of expressing bond strength in terms of bond order.
- Define the concept of bond length.
- Write down the resonance structure of the following NO2, SO3, and NO3.
- What is electronegativity? How does it differ from electron gain enthalpy?
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