Chemical Bonding Full Chapter | Class 11 Chemistry Chapter 4 | JEE 2025 Chemistry | Shilpi Mam

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The session focuses on JE Advanced 2024 Chemistry topics like Chemical Bonding, GOC, Basics of Organic Chemistry, and Coordination compounds in class 12th, with detailed explanations on bond formation, energy, and stability relationships. Key concepts like ionic, covalent, metallic, hydrogen bonds, and valence bond theories, along with hybridization, shape determination, and polarity are emphasized, highlighting the importance of understanding molecular structures and bond formations for exam preparation.

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The session focuses on JEE Advanced 2024 questions and key Chemistry topics like Chemical Bonding, GOC, and Basics of Organic Chemistry.
Detailed explanations are provided on the need for chemical bonds, with bond formation driven by attractive forces surpassing repulsive forces.
The session aims to complete vital chapters within a month, emphasizing the importance of understanding key concepts for exam readiness.
The relationship between energy and stability in chemistry is highlighted, with bond formation occurring at the state of minimum energy for stability.
Different types of strong bonds (ionic, covalent, metallic) exhibit varying bond energies (200-400 kJ/mol), requiring significant energy for breakage.
Weak bonds like hydrogen bonds and van der Waals forces have lower bond energies (2-40 kJ/mol), with hydrogen bonds being the weakest.
Understanding hybridization is crucial for predicting molecular geometry and bond formation, with specific geometries and angles determined by hybridized orbitals.

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Recent questions

What is the driving force behind bond formation?
Attractive force exceeding repulsive force.
How is bond energy related to stability?
Bond energy determines bond stability.
What is the Octet Rule in chemistry?
Atoms aim to have eight electrons in their outermost shell.
What are the types of strong bonds?
Ionic, covalent, and metallic bonds.
How are sigma and pi bonds formed?
Sigma bonds result from head-on overlapping, while pi bonds involve lateral overlapping.

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Summary

00:00

"JE Advanced 2024 Chemistry Preparation Session"

The session is focused on JE Advanced 2024 questions and important topics.
Chemistry topics like Chemical Bonding, GOC, and Basics of Organic Chemistry are discussed.
Coordination compounds will be covered in class 12th.
The session aims to complete important chapters within a month.
Detailed explanations on the need for chemical bonds and bond formation are provided.
The driving force behind bond formation is the attractive force exceeding repulsive force.
A graph illustrating potential energy and internuclear distance is analyzed.
Internuclear distance decreases as atoms come closer, leading to lower potential energy.
The relationship between energy and stability in chemistry is highlighted.
The session emphasizes understanding key concepts for better preparation and performance in exams.

14:50

"Energy, Stability, and Bond Formation in Chemistry"

The state with minimum energy is where bond formation occurs between hydrogen atoms at the closest distance.
Bond formation happens when the energy of the system is at its lowest point, indicating stability.
Bond energy is the energy at the lowest point from the origin, determining the stability of the bond.
Bond length indicates the internuclear distance, while energy signifies the bond's strength.
Repulsive forces dominate when atoms are brought closer beyond the point of bond formation.
The Octet Rule states that atoms aim to have eight electrons in their outermost shell for stability.
Hypovalent compounds occur when the central atom has fewer than eight electrons in its valence shell.
Hypervalent compounds involve central atoms with more than eight electrons in their valence shell.
Transition elements like zinc, cadmium, and mercury exhibit exceptions in electron configurations.
Odd electron species, such as NO2, demonstrate exceptions to the Octet Rule, forming coordinate bonds for stability.

30:23

Bond Types and Strengths in Chemistry

Coordinate bond involves sharing of electrons in a one-sided manner, with seven electrons in the valence shell, characteristic of odd electron species like ClO2.
Strong bonds, such as ionic and metallic bonds, exhibit interatomic interactions with bond energy ranging from 200 to 400 kilojoules per mole.
Strong bonds require significant energy to break, with energy released upon bond formation or absorbed during bond breakage.
Types of strong bonds include ionic, covalent, and metallic bonds, with bond energy ranging between 200 and 400 kilojoules per mole.
Weak bonds, like hydrogen bonds and van der Waals forces, have bond energy ranging from 2 to 40 kilojoules per mole, with hydrogen bonds being the weakest.
Classification of bonds based on strength includes strong bonds like ionic, covalent, and metallic bonds, and weak bonds like hydrogen bonds and van der Waals forces.
Kovalent bonds involve mutual sharing of electrons between two atoms, forming single, double, or triple bonds based on the number of shared electrons.
Electronegativity difference of less than 1.7 indicates covalent character, while differences above 1.7 suggest predominantly ionic character, varying based on different scales like the Pauling or Mulliken scale.
Polar covalent bonds result from unequal sharing of electrons due to differences in electronegativity, leading to partial charges on atoms and the introduction of polarity in the bond.
Non-polar covalent bonds, like in hydrogen molecules, occur when there is no electronegativity difference, resulting in no polarity or partial charges.

47:05

"Truth in Politics, Ionic and Coordinate Bonds"

Politics should not be introduced without knowing the truth.
Ionic bonds are formed by the transfer of valence electrons.
Sodium has one valence electron in the 3s orbital, while chlorine has 7 valence electrons.
Sodium loses electrons to achieve an octet, while chlorine gains electrons.
Ionic bonds form between atoms with a high electronegativity difference, typically a metal and a non-metal.
Coordinate bonds involve one atom contributing an electron pair while both atoms share the pair.
Coordinate bonds are also known as donor-acceptor or dative bonds.
Conditions for forming coordinate bonds include completing the octet of one atom and having a lone pair of electrons available.
Lewis dot structures represent elements by their symbol with dots or crosses indicating valence electrons.
When creating Lewis dot structures, identify the central atom based on its presence in the lowest number, ability to form multiple bonds, and highest electronegativity.

01:03:02

Identifying Central Atom and Lewis Dot Structures

Criteria for identifying the central atom in a molecule include making multiple bonds, being large-sized, and least electronegative.
The central atom in methane is carbon, which has four valence electrons.
Hydrogen atoms, each with one electron, share electrons with carbon to complete their duplet.
Through mutual sharing, carbon achieves an octet by gaining electrons from hydrogen.
The long-shot structure of methane shows carbon at the center with hydrogen atoms surrounding it.
Carbon monoxide's structure involves a coordinate bond between carbon and oxygen due to oxygen's octet completion.
The representation of a coordinate bond includes a plus sign on the donor atom and a minus sign on the acceptor.
In the carbonate ion, negative charges on oxygen atoms lead to single bonds with carbon and double bonds when neutral.
Resonance structures of the carbonate ion show the movement of double bonds and negative charges among oxygen atoms.
To determine the valence electrons for creating Lewis dot structures, subtract the charge (positive or negative) from the total valence electrons of the compound or ion.

01:18:59

Determining Electron Counts and Molecular Structures

The total number of valence electrons, n1, is determined to be 16 by calculating the valence electrons near nitrogen.
To find n2, the number of hydrogen atoms is divided by two, and the number of atoms excluding hydrogen is divided by eight.
The total number of electrons needed to complete everyone's octet is calculated to be 24.
N3 is calculated by halving n2, resulting in four bond pairs.
N4 is determined by subtracting n3 from n1, resulting in four lone pairs.
The structure of the Azide Ion is discussed, showing the completion of octets and formal charges.
The total number of electrons in the correct Lewis Octet structure of the Sulfite Ion is calculated to be 26.
The concept of hyper valent species is explained, with examples of molecules that do and do not follow the octet rule.
The importance of the Doubt App in clarifying concepts and providing solutions is highlighted.
The significance of understanding formal charges in determining molecular structures is emphasized.

01:35:55

Calculating Formal Charges and Valence Bond Theory

The formula for calculating formal charges is explained in the text.
Formal charges are determined by subtracting the actual valence electrons from the valence electrons in the Lewis dot structure.
The net charge on a polyatomic ion, such as carbonate, indicates the overall charge of the ion.
Formal charge is distinct from total net charge and is calculated individually for each atom in the ion.
The process of determining formal charges involves counting valence electrons and comparing them to the Lewis dot structure.
The formal charge on each atom is calculated by subtracting the actual valence electrons from the valence electrons in the structure.
The example of formal charge calculation is demonstrated for carbon and oxygen atoms in a carbonate ion.
The importance of understanding Valence Bond Theory (VBT) in explaining bond formation is highlighted.
VBT emphasizes the overlap of half-filled atomic orbitals with opposite spins for bond formation.
The extent and strength of bond formation depend on the nature of the orbitals and the quality of their overlapping, with non-directional orbitals resulting in weaker bonds compared to directional orbitals like p orbitals.

01:54:30

Orbital Overlapping and Bond Strength in Chemistry

Non-axial and axial types of orbitals are discussed, with d and s orbitals being directional and non-directional, respectively.
Directionality in orbitals leads to stronger bonds, with head-on overlapping being more effective than side wedge overlapping.
The extent of overlapping affects bond strength, with coaxial overlapping resulting in stronger bonds compared to lateral overlapping.
The nature of overlapping, whether head-on or side wedge, impacts the strength of the bond formed.
The value of n in orbitals influences bond strength, with larger n values leading to weaker bonds due to increased distance from the nucleus.
Carbon's superior catenation ability compared to silicon is attributed to the difference in bond strength based on n values.
The nature of orbital overlapping, whether PP, PS, or SS, determines the strength of the bond formed.
Positive, negative, and zero overlapping conditions are explained, with positive overlapping resulting in bond formation.
Proper orientation of atomic orbitals is crucial for successful overlapping and bond formation.
The type of overlapping, whether head-on or lateral, is determined by the molecular axis, with positive lateral overlapping being a specific type of overlap.

02:12:57

Formation of Sigma and Pi Bonds

Sigma bonds are formed from overlapping, with heads formed by axial overlap and side wedges formed by lateral overlapping.
Three combinations lead to sigma bonds: SS overlapping, S and P overlapping, and PP overlapping.
Sigma bonds are always formed by head-on overlapping along the internuclear axis.
The three cases for sigma bond formation involve S overlapping, S and P overlapping, and PP overlapping.
Sigma bonds are formed by overlapping S and P orbitals along the internuclear axis.
Pi bonds are formed by side wedge or lateral overlapping, with a half atomic number and opposite spin.
The formation of pi bonds involves overlapping from both sides, leading to bond shortening and increased bond energy.
Pi bond formation occurs by bringing P orbitals along different axes, resulting in pi bonds along the internuclear axis.
Pi bonds are formed by overlapping P orbitals along different axes, creating nodal planes where electron probability is zero.
Understanding sigma and pi bond formation is crucial for organic chemistry, as pi bonds have restricted rotation compared to sigma bonds.

02:31:23

Understanding Electron Pair Geometry and Molecular Shapes

The presence of electrons on Probability will be zero, indicating that in nodal plane, electron density will not be present.
To form a pie bond, P and D overlap, with the process involving taking P from the Axis and D from Aqua.
Overlapping P and D along the internuclear axis results in pie bond formation.
Side wedge overlapping is not suitable for forming pie bonds, with parallel or side wedge being the preferred method.
VBT explains the formation of sigma and pi bonds, emphasizing the importance of orientation and overlap.
The Vesper Theory focuses on valence shell electron pairs, distinguishing between bond pairs and lone pairs to determine molecular shape.
Repulsion between electron pairs influences molecular geometry, with lone pair-lone pair repulsion being the highest.
The steric number formula is used to determine the number of electron pairs near the central atom in a molecule.
Applying the steric number method to methane, with carbon having four valence electrons and forming four electron pairs with hydrogen atoms.
The steric number method aids in identifying the electron pairs near the central atom, crucial for determining molecular geometry and shapes.

02:48:41

Molecular Geometry and Bond Angles Explained

The molecule AB4 type has a central atom with no lone pair and a tetrahedral geometry.
The bond angles in this molecule are approximately 109.5 degrees.
For molecules with a central atom and no nearby lone pair, if there are two bond pairs, the geometry is linear with a bond angle of 180 degrees.
In molecules with three bond pairs, the geometry is trigonal with bond angles of 120 degrees.
For molecules with four bonded electron pairs, the common geometry is tetrahedral.
The steric number determines the geometry of a molecule based on the number of bond pairs and lone pairs.
In the case of steric number five, molecules with four bond pairs and one lone pair have a trigonal bipyramidal geometry.
The shape of molecules with a steric number of five can be altered by removing the lone pair, resulting in a seesaw shape.
Molecules with a steric number of four and three bond pairs, one lone pair have a trigonal pyramidal geometry.
Water, with a steric number of four and two bond pairs, two lone pairs, has a bent geometry.

03:05:52

Steric number determines molecular geometry and bonding

Steric number is crucial in determining the geometry of molecules
For steric number 6, with five bond pairs and one lone pair, the shape is octahedral
With five bond pairs and one lone pair, the shape is square pyramidal
For four bond pairs and two lone pairs, the shape is square planar
The position of the lone pair in square pyramidal shapes can vary without affecting the overall shape
In square planar shapes, the maximum distance between lone pairs is 180 degrees
The presence of lone pairs affects bond angles due to repulsion
Examples like SO2 and water demonstrate how lone pairs impact bond angles
Identical shapes in molecules can be determined by the presence of lone pairs and bond pairs
Understanding hybridization is essential for predicting molecular geometry and bond formation

03:22:49

Creating Hybrid Orbitals for Stable Bonds

The concept involves mixing atomic orbitals to create new hybrid orbitals for energy conservation.
By intermixing 2s and 2p atomic orbitals, four new hybrid orbitals are formed.
Hybrid orbitals are similar in shape, size, and energy, ensuring stable bonds.
Different types of hybridized compounds are formed, such as sp2, sp3, sp3d, sp3d2, and sp3d3.
Hybrid orbitals have specific geometries, like linear, trigonal planar, and tetrahedral.
The percentage of s character in hybrid orbitals can be calculated based on the number of s and p orbitals involved.
Exciting electrons in carbon atoms can lead to the formation of sp3 hybridized orbitals for bonding.
The geometry of hybridized orbitals is crucial for understanding their properties and interactions.
Steric number helps determine the type of hybridization, such as sp3, sp3d, sp3d2, or sp3d3.
The arrangement of bonds in hybridized compounds, like in PClF, follows specific angles and positions for stability.

03:40:03

Hybridization and Geometry in Chemical Compounds

120-degree angles are present in the geometry, with three such angles formed by chlorine and phosphorus atoms.
Axial bonds in the geometry are longer and weaker compared to equatorial bonds to minimize repulsion.
Heating PCL5 at high temperatures causes the axial chlorine atoms to break away, leaving behind PCL3.
Calculating hybridization involves determining the sum of valence electrons and dividing by specific numbers based on the total.
The steric number is crucial in determining hybridization, with specific calculations leading to different hybridization types.
Methane's hybridization is determined by dividing the total valence electrons by two, resulting in an sp3 hybridization.
Valence electron calculations for sulfur tetrafluoride (SF4) lead to an sp3d hybridization due to the steric number.
PCL5 exists in both sp3 and sp3d2 hybridizations in the solid state, with PCL4 existing in an ionic form.
Bains Rule guides the placement of lone pairs and multiple bonds in molecular geometry based on electronegativity.
Understanding hybridization in odd-electron molecules involves considering the side atom's electronegativity and the presence of an odd electron.

03:57:49

Electronegativity Impact on Hybridization and Bond Angles

Electronegativity determines the treatment of electrons as lone pairs and impacts hybridization.
More electronegative atoms are treated as lone pairs, affecting the steric number and hybridization.
The electronegativity order is Fluorine, Oxygen, Nitrogen, and Carbon.
Electrons participate if the side atom is more electronegative, influencing hybridization.
Bond angles are compared based on hybridization, with higher s character resulting in larger bond angles.
The presence of lone pairs on the central atom decreases bond angles due to repulsion.
Electronegativity of the central atom directly affects bond angles, inversely related to the side atom's electronegativity.
The Dapor moment measures polarity in molecules, with the arrow pointing towards the more electronegative atom.
Dapor moment is zero in molecules with regular geometry and when side atoms are the same.

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