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Chemical Bonding FULL CHAPTER | Class 11th Inorganic Chemistry | Arjuna NEET

Arjuna NEET・170 minutes read

Chemical bonding in organic and inorganic chemistry is crucial for understanding how atoms form molecules through various types of bonds like ionic and covalent, with factors like VSEPR theory, hybridization, and energy playing key roles in bond formation and stability. The text also delves into concepts like resonance, dipole moments, and hybrid orbitals, emphasizing the importance of understanding these principles in determining molecular geometry and bond strength.

Insights

Understanding chemical bonding is crucial for maintaining friendships between atoms.
Different types of bonding, including ionic and covalent, are explored in detail, along with resonance and dipole moments.
The chapter covers various theories like VSEPR theory, hybridization, and bond length, emphasizing the importance of understanding hybridization before delving into other topics.
Ionic bonding involves the transfer of electrons between metals and non-metals to form oppositely charged ions, leading to the creation of ionic compounds with specific properties.
The concept of resonance plays a significant role in understanding molecular structures, with resonance energy, stability, and the number of resonating structures impacting the overall stability of a compound.

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Recent questions

What is the difference between ionic and covalent bonds?
Ionic bonds involve electron transfer, while covalent bonds share electrons.

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Summary

00:00

Chemical Bonding: Essential Concepts and Theories

Organic chemistry is divided into chapters, with chemical bonding being a crucial one.
Understanding chemical bonding is essential for maintaining friendships between atoms.
Inorganic chemistry's base is established in the last chapter, focusing on trends and bonding.
The chapter delves into how atoms combine to form molecules through bonding.
Different types of bonding, such as ionic and covalent, are discussed, along with resonance and dipole moments.
The chapter covers theories like VSEPR theory, hybridization, bond length, angle, and energy.
Hydrogen bonding and MOT are also explored, along with previous year questions.
The importance of understanding hybridization before proceeding with other topics is emphasized.
Chemical bonding involves forces of attraction between atoms and ions in molecules.
Ionic bonding involves the transfer of electrons between metals and non-metals to form oppositely charged ions.

19:15

Formation and Properties of Ionic Compounds

Metal does not lose electrons; non-metal generates the electron.
Ionic closure forms between metal and non-metal, creating an ionic bond.
Sodium chloride formation involves sodium losing an electron and chlorine gaining one.
The configuration of sodium is 1s2 2s2 2p6 3s1, while chlorine's is 1s2 2s2 2p6 3s2 3p5.
The formation of oppositely charged ions leads to ionic closure.
Lewis dot structure helps represent ionic bonds, showing electron transfers.
Calcium loses two electrons, while oxygen gains two, forming a neutral compound.
Magnesium chloride and sodium sulfide formation involve electron transfers between metals and non-metals.
Factors affecting ionic bond formation include metal's ionization potential, non-metal's electron affinity, and lattice energy.
Ionic compounds have high melting and boiling points, are hard and brittle, and do not conduct electricity in solid form due to fixed ions.

40:51

Ionic and Covalent Bonds: Conductivity and Structure

Ions in compounds are fixed in their positions but conduct electricity when melted or dissolved in water.
Electrovalent compounds in the molten state dissociate into ions, allowing them to conduct electricity.
Ionic compounds are soluble in polar solvents like water due to their polarity.
Electrovalent bonds are non-directional and not fixed, leading to fast ionic reactions when dissolved in water.
Isomorphism occurs in compounds with similar chemical compositions and crystalline structures.
Lattice energy is higher when the product of charge is greater and the distance between ions is smaller.
Lattice energy is lower when the distance between ions is greater, as seen in larger-sized ions.
Covalent bonds involve the mutual sharing of electrons between non-metal atoms, forming directional bonds.
Covalent bonds are formed between non-metal atoms, leading to stable electron configurations.
Examples of covalent bonds include H2, N2, and H2O, where electrons are shared to complete octets or doublets.

01:04:48

Covalent Bonds: Structure, Classification, and Properties

Closed feet, three closed legs on loan, one on loan
Methane structure: carbon with four valence electrons, hydrogen sharing electrons with carbon
Carbon dioxide structure: carbon sharing electrons with oxygen to complete octet
Classification of elements into covalent and ionic based on metal or non-metal presence
Classification of valence bonds into single, double, and triple based on electron sharing
Electronegativity and polarity in covalent bonds: chlorine attracting electrons more than hydrogen
Examples of non-polar covalent bonds: F2, Cl2, O2, N2
Examples of polar covalent bonds: water, ammonia
Properties of covalent compounds: mostly gases or liquids, some exist as soft solids, high melting points in 3D network structures
Solubility of covalent compounds: soluble in polar solvents, some soluble in non-polar solvents due to hydrogen bonding

01:26:38

Electron Sharing and Lewis Acids in Chemistry

Sharing of electrons between A and B results in two electrons being shared.
The electron sharing is unidirectional, with A sharing its two electrons with B.
The coordination bond formed between A and B behaves like a covalent bond.
The compound properties after coordination resemble those of valence compounds.
The structure of ammonia involves nitrogen donating electrons to hydrogen, resulting in a positively charged nitrogen.
Oxygen in water structure has two lone pairs of electrons, leading to electron sharing with hydrogen.
Boron trifluoride (BF3) is an electron-deficient compound, requiring electron donation from ammonia.
Lewis acids, like BF3, accept electrons from ammonia to stabilize their electron deficiency.
Ozone formation involves the splitting of O2 molecules to create O3 molecules.
Calculating formal charges on atoms involves subtracting lone pair and bonding electrons from valence electrons.

01:54:33

"Resonance in Molecules: Fluctuating Bond Types"

Phosphorus in PCl3 has five balanced electrons, with three chlorines each gaining one electron, totaling eight electrons.
In CCl4, carbon and four chlorines achieve octet completion, with carbon sharing its electrons with the chlorines.
The octet rule is not always followed, as seen in compounds like SF6, where sulfur has 12 electrons.
Noble gases typically do not form compounds due to their complete octets, except for some exceptions like xenon fluorides.
The concept of resonance hybrid is illustrated through examples like ozone and benzene, where structures fluctuate between single and double bonds.
In ozone, the resonance hybrid shows equal bond lengths due to the fluctuation between single and double bonds in the oxygen atoms.
Benzene also exhibits resonance, with carbon atoms alternating between single and double bonds, resulting in equal bond lengths.
The resonance hybrid of carbonate ions showcases the fluctuation between single and double bonds in the oxygen atoms, leading to equal bond lengths.
The Resonating Structures or Canonical Forms contribute to the Resonance Hybrid, representing the actual structure of the molecule.
The Resonance phenomenon is observed through various examples, demonstrating the concept of fluctuating bond types within molecules.

02:17:47

Understanding Resonance and Dipole Moments

Resonance structures are explained, with the resonance hybrid being the actual structure.
The stability of resonating compounds is high, with the goal of creating as many resonating structures as possible.
The stability of a compound increases with the number of resonating structures it has.
Resonance energy is calculated by subtracting the energy of the most stable resonating structure from the energy of the resonance hybrid.
The total number of bonds between two atoms in a structure is divided by the total number of resonating structures to calculate the closed order.
Examples of calculating resonating structures are provided for carbonate ion and phosphate ion.
Dipole moment is defined as the product of the magnitude of charges and the distance between them.
The unit of dipole moment is represented as 1*10^-18 esu in centimeters.
The presence of charge separation in a molecule determines its polarity and dipole moment.
The dipole moment in diatomic molecules is zero for homonuclear atoms, while in polar molecules, the dipole moment is non-zero due to electronegativity differences.

02:39:48

Molecular Polarity and Ionic Character in Bonds

BF3 molecule is non-polar due to the cancellation of dipole moments.
Ammonia is a polar molecule with a non-zero dipole moment.
Water is a polar molecule with a higher dipole moment than hydrogen sulfide.
Methane is a non-polar molecule with a dipole moment of zero.
The dipole moment of a molecule decreases with decreasing electronegativity difference.
The dipole moment of H2 molecule is zero, making it non-polar.
The dipole moment of chloroform indicates its polarity.
The percentage ionic character in a covalent bond can be calculated using the equation.
Electronegativity values determine the order of elements in a covalent bond.
Ionic compounds insoluble in water may indicate the presence of ionic character in covalent compounds.

03:02:55

Ionic Bonding and Molecular Geometry Explained

Protons and electrons are discussed in relation to different elements, such as magnesium and aluminum.
Anions have more electrons than protons, leading to attraction and distortion of electron clouds.
The concept of polarizing power is introduced, affecting the electron cloud of ions.
The size and charge of ions determine their polarizing power and covalent character.
The electronic configuration of ions like Copper Plus influences their polarizing power.
The theory of valence electron pair repulsion explains the geometry of molecules.
Lone pairs of electrons experience the most repulsion, followed by lone pairs participating in bonding.
The order of repulsion intensity is lone-lone, lone-bonding, and bonding-bonding pairs.
The number of electron pairs in a molecule determines its geometry, such as linear or trigonal planar.
The angle between bonds in a molecule is determined by its geometry, like 180° for linear and 120° for trigonal planar.

03:26:04

Molecular Geometry: Angles and Lone Pairs

1009.5° is rounded off from 1009.5° or 100 90 degrees, with 28 minutes being approximately 0.5 degrees.
A molecule with five electron pairs exhibits trigonal pyramidal geometry.
In trigonal pyramidal geometry, there are three atoms in the plane, forming equatorial positions.
Trigonal planar geometry consists of two types of bond angles: 120° and 90°.
The closed angle between two atoms in trigonal planar geometry is 120°.
In octahedral geometry, there are five closed angles and one lone angle, with a 90° closed angle.
Pentagon bipyramidal geometry features a closed angle of 72° and 90°, with 180° in the middle.
Octahedral geometry shifts to square pyramidal when one lone angle is introduced.
The presence of lone angles in molecules can lead to geometry distortion due to repulsion.
Calculating the total valence electrons in a molecule involves dividing the total balance by eight to determine the number of lone pairs and achieve a final reminder of zero.

03:49:41

Balancing Electron Distribution in Molecules

Repulsion is a key concept in balancing electron distribution.
The example of SF6 is used to illustrate the calculation of total balance electrons.
Fluorine has 42 valence electrons, sulfur has 6.
Dividing 42 by 8 gives 5, with a remainder of 2.
The geometry of SF6 is octahedral.
The example of SF4 is quickly analyzed.
The total balance electrons in SF4 are calculated to be 34.
The geometry of SF4 is trigonal bipyramidal.
The rules for determining geometry based on electronegativity are explained.
The example of PCl3F2 is used to demonstrate the application of these rules.

04:20:47

Understanding Molecular Geometry and Bond Formation

The text discusses the concept of molecular geometry and the calculation of central atoms.
It mentions the drawback of the VSEPR theory in explaining molecular shapes.
The text introduces the concept of orbital overlap and its importance in covalent bond formation.
It explains that electrons in overlapping orbitals should have opposite spins for stability.
The strength of a bond is directly related to the extent of orbital overlap.
The text emphasizes that greater overlap leads to stronger and more stable bonds.
It explains that shorter bond lengths indicate stronger bonds due to increased overlap.
The stability of a bond is determined by the energy required to break it.
The text discusses the classification of covalent bonds into sigma and pi bonds based on orbital overlap.
It explains the process of hybridization in forming new sets of orbitals with similar energies.

04:48:11

Hybridization Shapes Molecules with Distinct Geometry

Hybridization involves mixing atomic orbitals of similar energy to form new hybrid orbitals.
Hybridization occurs when atomic orbitals mix to form new hybrid orbitals with distinct geometries.
The geometry of hybrid orbitals is crucial in determining the shape of molecules.
Hybrid orbitals are formed to minimize electron repulsion and achieve stable molecular geometry.
Different types of hybridization, such as sp, sp2, and sp3, are observed in carbon atoms.
The hybridization of carbon atoms is determined by the number and type of sigma and pi bonds they form.
The hybridization of carbon atoms in benzene results in sp2 hybridization due to the three sigma and one pi bond configuration.
The bond length between atoms is influenced by factors such as the size of the atoms involved.
Ionic bonds affect bond length by adding the radii of the ions involved.
The bond length is directly proportional to the size of the atoms forming the bond.

05:20:29

Caste and Atom Size Influence Bond Length

Increasing caste leads to an increase in bond length, as seen in HDFC.
The bond length is influenced by the size of the atom, with longer bond lengths indicating larger atoms.
Energy required to break a bond is directly proportional to bond strength and inversely proportional to bond length.
Exceptions exist in halogens, where chlorine requires less energy to break the bond compared to other halogens due to repulsion factors.

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