Chemical Bonding and Molecular Structure

Chemical Bonding and Molecular Structure

• Structure dictates chemical properties in compounds
• The strength and directionality of inter-atomic forces are key factors.

Chemical Bonds

• This refers to the attractive force between two atoms that binds them together to form a compound or molecule
• Stability in nature correlates with bond formation
• Elements seek inert electronic configurations for maximum stability
• Noble gas electron configuration can be achieved through transference, mutual sharing, or donation of lone pair electrons.

Types of Bonds

• Ionic bonds, Covalent bonds, Coordinate bonds Metallic bonds, Hydrogen bonds, Van der Waals bonds
• Formation of chemical bonds can be explained in terms of electrons using Lewis structures.
• Atoms form stable octets when linked by a chemical bond, as per Lewis's postulation

Lewis Dot Structures

• In molecule formation, valence electrons are critical
• Lewis symbols represent an element with its symbol and valence electrons

Octet Rule

• Every atom aims to achieve a noble gas electronic configuration, typically having eight valence electrons and hence referred to as the law of octet.
• The 'law of duplet' applies if an atom needs to have two valence electrons
• Only stable compounds adheres to the octet rule in Lewis's view

Formal Charge

• Formal charge indicates the difference between valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure.
• Formal Charge = (Number of valence electrons in neutral atom) - (Number of unshared electrons + 1/2 Number of shared electrons)
• Bonded atoms vs Unbonded atoms calculation

Ionic Bond

• Via complete electron transfer forms stable noble gas electronic configurations
• The atom losing electrons becomes a positive ion and the one gaining becomes a negative ion.

Ionic Bonding

• Electrovalency refers to the number of electrons lost or gained to form an ionic or electrovalent bond.
• Ionic compounds tend to be hard, brittle, and crystalline
• They exhibits high melting and boiling points
• Ionic compounds are polar but linkage between oppositely charged ions is non-rigid and non-directional.
• Soluble in polar solvents like water, yet insoluble in non-polar solvents such as CCl4, or benzene
• Electricity is conducted well in fused or solution states but poorly in solid states due to ion mobility

Covalent Bond

• Mutual electron sharing between atoms forms a covalent bond.
• Identical atoms yields homoatomic molecules, while different atoms yields heteroatomic molecules.

Valence Bond Theory (VBT)

• Heitler and London introduced the valence bond theory in 1927 later developed by Pauling
• Electronic configurations and atomic orbitals underlie VBT
• Attractive and repulsive forces operate when two hydrogen atoms (A & B with nuclei NA & NB, electrons eA & eB) approach and combine.
• Nucleus of one atom attracts its own electron and the electron of the other atom, and vice-versa.
• Repulsive forces exist between electrons of two atoms and nuclei of two atoms.
• Closer atoms is due to attractive forces and repulsive forces make them move apart.

Orbital Overlap Concept

• Atomic orbitals must undergo partial interpenetration to minimize energy during hydrogen molecule formation.
• Partial interpenetration of atomic orbitals forms overlapping of atomic orbitals
• Overlap between atomic orbitals can be positive, negative, or zero based on the orbitals.

Types of Overlapping

• Covalent bonds are classified into two categories based on overlapping type.
• A sigma (σ) bond forms through axial overlapping of half-filled atomic orbitals along the internuclear axis.
• S-S overlap: Two half-filled s-orbitals overlap along the internuclear axis.
• S-P overlapping: A half filled s-orbital of one atom overlaps with a half filled p-orbitals of the other atom, which further forms a s-p sigma bond.
• P-P overlapping: Co-axial overlapping occurs between half filled p-orbitals of two atoms, creating p-p sigma bond.
• Pi (π) Bond: Atomic orbitals align parallel and perpendicular to the internuclear axis when covalent bonds form
• Sidewise overlapping forms two saucer-shaped charged clouds above and below the participating atoms.
• Electrons in the π bond formation are called pi-electrons.

Hybridisation

• Hybridisation is intermixing of orbitals with slightly different energies, resulting in new orbitals with equivalent energies and shape
• Atomic orbitals combine to form hybrid orbitals.

Salient Features of Hybridization

• Hybrid orbitals equals the number of atomic orbitals that get hybridized

• Hybridized orbitals always has equivalent energy and shape.

• Hybrid orbitals are more effective and stronger than atomic orbitals.

• Molecule's geometry depends on the type of hybridisation

Important conditions for hybridisation:

• Orbitals in the valence shell of the atom are hybridised.
• There should be small differences in the energies of hybridized orbitals.
• Electron promotion is not essential for hybridisation.
• Half-filled orbitals don't need to participate in hybridisation.

Types of Hybridisation

• Mixing depends on the type of orbitals involved, like sp³, sp², sp, sp³d, or sp³d².
• Sp-hybridisation: one s and one p orbitals mix to produce two equivalent hybrid orbitals, known as sp hybrid orbitals.
• The two sp-hybrid orbitals are arranged in a straight line, creating a 180° angle and the molecule has linear geometry, so the hybrid orbitals has 50% s-character and 50% p-character.
• Molecules with sp-hybridisation are BeF2, BeCl2, and BeH2
• Sp²-hybridisation: mixing of one s and two p orbitals produce three equivalent hybrid orbitals, known as sp² hybrid orbitals.
• Sp² hybrid orbitals are larger in size than sp-hybrid orbitals but slightly smaller than that of sp³ hybrid orbitals.
• Each sp² hybrid orbitals possess equal to 1/3 (or 33.33%) s-character and 2/3 (or 66.7%) p-character, molecules of with are BF3, BCl3, or BH3
• Sp³d-hybridisation: Mixing one s, three p, and one d-orbitals produce five sp³d hybridised orbitals, adopting a trigonal bipyramidal shape.

Formation of PCl5

• Phosphorus ground state electronic configuration is 1s² 2s² 2p⁶ 3s² 3p³
• Under certain conditions, the 3s-electrons can unpair, and one electron moves to the vacant 3dz² orbital.
• Ground and excited state configurations of phosphorus are shown with electrons contributed by Cl atoms

Sp³d²-hybridisation

• One s, three p and two d-orbitals form six identical sp³d² hybrid orbitals to undergo intermixing
• The six orbitals are oriented towards the corners of an octahedron, in space at a 90° angle to each other.
• Sulfur has a ground state outer configuration of 3s² 3p⁴, whereas it excited state is when electrons pairs get unpaired in 3s and 3px orbitals and one pair is promoted to 3dz² and 3dx²-y² orbitals.

Valence Shell Electron Pair Repulsion (VSEPR) Theory

• Sidgwick and Powell introduced it in 1940 and was further developed by Nyholm and Gillespie
• Key postulates depends on the number of electron pairs around the central atoms the molecule's shape.
• The electron pairs seek to minimize repulsion around the central atom, as the valence shell sphere is negatively charged.
• The valence shell is treated as a sphere with electron pairs at maximum distance
• Multiple bonds behave as a single electron pair.

Bond Parameters

• Bond Angle: The angle between orbitals containing bonding electron pairs around the central atom, typically measured in degrees via spectroscopic methods
• Bond Length: Equilibrium distance between the nuclei of two bonded atoms
• Lattice Enthalpy: The energy it takes to break apart one mole of a solid ionic compound to gaseous constituent ions
• NaCl example is 788 kJ mol-1
• Bond Order: Half the difference between number of electrons in bonding and antibonding molecular orbitals, is a whole, fraction, or zero and is positive or negative.
• Bond Enthalpy: Energy required to break one mole of bonds of a particular type between gaseous atoms, with units in kJ mol-1. Hydrogen molecule H – H bond is 435.8 kJ mol-1.

Molecular Orbital Theory (MOT)

• Hund and Mulliken developed it, and this theory considers molecules differently from their constituent atoms.
• All electrons in the atoms form the constitution of a molecule move in the entire molecule
• Molecules possess varying energy level orbitals like atoms, known as molecular orbitals.

Energy Level Diagram for Molecular Orbitals

• Antibonding orbitals exhibit increased energy than atomic orbitals

Resonance

• When right frequency excites a metal, then photo electric effect occurs
• Compounds with same molecular formula, differing structural formulas, are known as resonating or canonical structures
• None structure explains all properties.

Hydrogen Bonding

• Hydrogen attached to highly electronegative elements (nitrogen, oxygen, fluorine) creates a covalent bond
• Covalent bond's electrons get pulled to the electronegative atom
• Forms hydrogen bonds and is weaker than a covalent bond
• Two types of hydrogen bonding depends on if it is similar or dissimilar molecules
• Intramolecular hydrogen bonding takes place within the molecules
• Affects compound conditions for H2O is liquid at room temperature but H2S is gas due to intermolecular hydrogen bonds in H2O
• Only covalent molecules that tend to form intermolecular hydrogen bonds are water-soluble

Metallic Bonding

• Force binding a metal atom and number of electrons
• Electrical, thermal conductivity, bright metallic luster, malleability, ductility, tensile strength, elasticity are impacted by this bonding

Bond Characteristics

• When atoms get closer, attraction occurs with decreasing potential energy, until it becomes minimum, and when the atoms get further, repulsion starts and system starts to increase.
• Equilibrium distance and atom vibration happens at mean position and center of nucleis of two bonded atoms is bond length, where it is called bond length
• Express in terms of angstrom or picometer.
• Determined experimentally by x-ray diffraction, electron diffraction, or spectroscopic methods.
• Bond length in ionic compounds is the sum of ionic radii as well as covalent compounds is the sum of covalent radii.
• For a covalent molecule AB, the bond length is given by d= ra + rb

Factors affecting Bond length

• Size of atoms, Bond length increases if there is increase in atom size
• Multiplicity of Bond, Bond length reduces with stronger bond
• Type of hybridisation, S orbital is smaller in size, greater the s character yields shorter hybrid orbitals which causes a shorter bond.

Bond enthalpy

• When atoms bind together, energy gets released
• Breaking one mole of bonds requires separating them into gaseous atoms which is called bond dissociation enthalpy or Bond enthalpy and is measured in KJ mol-1
• Greater is the bond dissociation enthalpy,stronger is bond.
• Diatomic molecules H2, Cl2, O2, N2, HCl, HBr, Hl the bond enthalpies equals to their dissociation enthalpy and is fixed.
• Polyatomic molecules does not posse the same bond

Factors Affecting Bond Enthalpy

• Size decrease in the atom increases bond dissociation enthalpy and bond strength
• Greater are the multiplicity of bonds, the bond dissociation enthalpy is also greater.
• The repulsion between atoms is greater when there is too much electrons, and the bond dissociation enthalpy decreases.
• Bond Angles: Form when overlap is forming or breaking a bond
• The angle on which electrons are binding is called Angles

Bond Order

• Bond order, greater is bond enthalpies and stability and greater the bond order, the shorter is bond length.

Polar and Non-Polar Covalent Bond

• In non polar molecule both atoms are attracted equally hence there is no poles are developed, where as for hydrogen both electron lie in middle.
• The electron cloud is symmetrical and there is no charge separation.
• For polar bond, atoms attract more electrons, hence the electron cloud is displacing electrons that makes bond polar bond.

Dipole Moments

• Each polar molecule has two poles where the molecule possesses an electric field.
• Negative charges is always equal to the positive charge, and the product of multiplying them with center is dipole moment
• The charge q is of the order of 10-10 esu and the internuclear distance d is of the order of 10-8 cm.
• In S.I. units, 1 D = 3.335 × 10 -30 Cm
• Dipole moment is a vector with direction
• If molecule is diatomic, then both molecule is same.

Fajan's Rules

• Anion approaches cations to attracted to it and bends it which is polarization.
• The polarisability is tendency an anion gets polarization is the polarization of is strength anion

Factors On Polarizing Power

• Smaller cation
• Larger anion
• More change can cause polarized polarization, greater power polarizes power
• Less easily be easily polarized

Characteristics of covalent compounds

• Can be physical. such as existing solids, liquids, gases
• There crystal molecule
• High levels need for electrical conductivity in the conduct
• Rigid structures form structural Isomers

Shapes Of Molecules

• Shapes for a few molecules
• Beryllium fluoride atomic is 4 is ²2s², In excited which oriented at an angle pf 180,
• Boron trifluoride the molecule is 5, which are a and is oriented at degree to
• Shape of methene are Tetrahedral in shapes to which are arranged at 109

Shapes of Ethane molecule

• Each has its is sp³ so they have 109
• They of is thus the bonding of internucleated each atom's sphere

Shape for Ethylene

• In is the sp ². the atoms are linear to one of p,z

Shape of Actylene molecules

• Each is of Sp, one is Sp so both form electron in a SP that linear to a C axis both axes are in back of it

Key terminlogy:

• chemical bond forms when there are multiple

• electrovalent bonding, is is when electrons are changed electrons forming is electrovalency

• Covalency is filled bonds for is

• The shared bonding is Dative

• Hybrid formation of energy

• Is geometrical arrangement when

• Regular geometrical is possess geometrical, If the repulsion of interaction, then there as irregular.

• Electronegativity power Is to attract in bonding

• Bonding Dipole is when the of moment is and

• Are the if or are

• The of force of in molecules.

• Hydrogen in one can bonded force.

• The is orbital structure or as hybrids

• The is linear combined of of orbitals.