Chemical Bonding and Molecular Structure

Introduction

Matter is composed of elements, which are substances made up of only one type of atom. Under normal conditions, no element exists as independent atoms except noble gases. Instead, atoms combine to form groups known as molecules, which have characteristic properties different from their constituent atoms. These molecules are formed due to chemical bonding, which is the force of attraction holding atoms together. Understanding chemical bonding is crucial because it explains the structure, properties, and behavior of substances. The study of chemical bonding involves exploring how atoms achieve stable electronic configurations, often resembling the nearest noble gas configuration, which is generally more stable. This stability is often explained by the octet rule, which states that atoms tend to have eight electrons in their valence shell to attain stability. However, there are exceptions and limitations to this rule. The chapter introduces the Kössel-Lewis approach to chemical bonding, which provides a simplified way to visualize bonding through electron transfer or sharing to complete octets. This foundational knowledge sets the stage for understanding different types of bonds, molecular structures, and theories explaining bonding at the atomic and molecular levels. **Table on page 1 (1×2)** | MOLECULAR STRUCTURE Scientists are constantly discovering new compounds orderly arranging the facts about them, trying to explai with the existing knowledge, organising to modify th earlier views or evolve theories for explaining the newl After studying this Unit, you will be observed facts. able to • understand Kössel-Lewis approach to chemical bonding; • explain the octet rule and its Matter is made up of one or different type of elements limitations, draw Lewis structures Under normal conditions no other element exists as a of simple molecules; independent atom in nature, except noble gases. However a group of atoms is found to exist together as one specie • explain the formation of different types of bonds; having characteristic properties. Such a group of atom is called a molecule. Obviously there must be some forc • describe the VSEPR theory and which holds these constituent atoms together in th predict the geometry of simple molecules. The attractive force which holds variou molecules; constituents (atoms, ions, etc.) together in differen • explain the valence bond chemical species is called a chemical bond. Since th approach for the formation of formation of chemical compounds takes place as a result o covalent bonds; combination of atoms of various elements in different ways • predict the directional properties it raises many questions. Why do atoms combine? Why ar of covalent bonds; only certain combinations possible? Why do some atom combine while certain others do not? Why do molecule • explain the different types of possess defni ite shapes? To answer such questions differen hybridisation involving s, p and d orbitals and draw shapes of theories and concepts have been put forward from tim simple covalent molecules; to time. These are Kössel-Lewis approach, Valence She Electron Pair Repulsion (VSEPR) Theory, Valence Bond (VB • describe the molecular orbital Theory and Molecular Orbital (MO) Theory. The evolutio theory of homonuclear diatomic molecules; of various theories of valence and the interpretation o the nature of chemical bonds have closely been related t • explain the concept of hydrogen the developments in the understanding of the structur bond. | | | --- | --- | | | , n e y . n , s s e e s t e f , e s s t e ll ) n f o e | **Table on page 8 (4×3)** | energy absorbed. Thus a qualitative measure of the stability of an ionic compound is provided by its enthalpy of lattice formation and not simply by achieving octet of electrons around the ionic species in gaseous state. Since lattice enthalpy plays a key role in the formation of ionic compounds, it is | | | | --- | --- | --- | | important that we learn more about it. 4.2.1 Lattice Enthalpy The Lattice Enthalpy of an ionic solid is defined as the energy required to completely separate one mole of a solid ionic compound into gaseous constituent ions. For example, the lattice enthalpy of NaCl is 788 kJ mol–1. This means that 788 Fig. 4.1 The bond length in a covalent kJ of energy is required to separate one mole molecule AB. of solid NaCl into one mole of Na+ (g) and one R = r + r (R is the bond length and r and r ar A B A B mole of Cl– (g) to an infinite distance. the covalent radii of atoms A and B respectively) This process involves both the attractive forces between ions of opposite charges in the same molecule. The van der Waal and the repulsive forces between ions of radius represents the overall size of th like charge. The solid crystal being three- atom which includes its valence shell in dimensional; it is not possible to calculate nonbonded situation. Further, the van de lattice enthalpy directly from the interaction Waals radius is half of the distance betwee of forces of attraction and repulsion only. two similar atoms in separate molecules i Factors associated with the crystal geometry a solid. Covalent and van der Waals radii o have to be included. chlorine are depicted in Fig. 4.2. 4.3 Bond Parameters r = 99 pm 4.3.1 Bond Length c 198 pm Bond length is defined as the equilibrium distance between the nuclei of two bonded atoms in a molecule. Bond lengths are measured by spectroscopic, X-ray diffraction and electron-diffraction techniques about which you will learn in higher classes. Each rvdw atom of the bonded pair contributes to the = 180 bond length (Fig. 4.1). In the case of a covalent bond, the contribution from each atom is pmmp called the covalent radius of that atom. 360 The covalent radius is measured approximately as the radius of an atom’s Fig. 4.2 Covalent and van der Waals radii i | | | | | | e s e a r n n f | | | r = 99 pm 198 c pm rvdw = 180 pmmp 360 | | | | | n | **Table on page 9 (8×4)** | Table 4.2. Bond lengths for some common molecules are given in Table 4.3. The covalent radii of some common elements are listed in Table 4.4. 4.3.2 Bond Angle | Bond Type | Covalent Bond Length (pm) | | | --- | --- | --- | --- | | | O–H C–H N–O C–O | 96 107 136 143 143 154 121 122 133 138 116 120 | | | It is defined as the angle between the orbitals C–N 143 containing bonding electron pairs around the C–C 154 central atom in a molecule/complex ion. Bond C=O 121 angle is expressed in degree which can be N=O 122 experimentally determined by spectroscopic C=C 133 methods. It gives some idea regarding the C=N 138 distribution of orbitals around the central C≡N 116 atom in a molecule/complex ion and hence it C≡C 120 helps us in determining its shape. For Table 4.3 Bond Lengths in Some Commo example H–O–H bond angle in water can be Molecules represented as under : Molecule Bond Length (pm) H (H – H) 74 2 F (F – F) 144 2 4.3.3 Bond Enthalpy Cl (Cl – Cl) 199 2 It is defined as the amount of energy required Br (Br – Br) 228 2 to break one mole of bonds of a particular I (I – I) 267 2 type between two atoms in a gaseous state. N (N ≡ N) 109 2 The unit of bond enthalpy is kJ mol–1. For O (O = O) 121 2 example, the H – H bond enthalpy in hydrogen HF (H – F) 92 molecule is 435.8 kJ mol–1. HCl (H – Cl) 127 HBr (H – Br) 141 H (g) → H(g) + H(g); ∆ H = 435.8 kJ mol–1 2 a HI (H – I) 160 Similarly the bond enthalpy for molecules containing multiple bonds, for example O and Table 4.4 Covalent Radii, *r /(pm) 2 cov N will be as under : 2 O (O = O) (g) → O(g) + O(g); 2 ∆ H = 498 kJ mol–1 a N (N ≡ N) (g) → N(g) + N(g); 2 ∆ H = 946.0 kJ mol–1 a It is important that larger the bond dissociation enthalpy, stronger will be the bond in the molecule. For a heteronuclear diatomic molecules like HCl, we have HCl (g) → H(g) + Cl (g); ∆ H = 431.0 kJ mol–1 a In case of polyatomic molecules, the | C–N C–C C=O N=O C=C C=N C≡N C≡C | 143 154 121 122 133 138 116 120 | | | | | | n | | | Molecule | Bond Length (pm) | | | | H (H – H) 2 F (F – F) 2 Cl (Cl – Cl) 2 Br (Br – Br) 2 I (I – I) 2 N (N ≡ N) 2 O (O = O) 2 HF (H – F) HCl (H – Cl) HBr (H – Br) HI (H – I) | 74 144 199 228 267 109 121 92 127 141 160 | | | | | | | | | | | | | measurement of bond strength is more | | | | **Table on page 13 (6×6)** | dipole because of lone pair decreases the effect (n-1)dnnso, typical of transition metals, i of the resultant N – F bond moments, which more polarising than the one with a nobl results in the low dipole moment of NF as 3 gas confgi uration, ns2 np6, typical of alka represented below : and alkaline earth metal cations. The cation polarises the anion, pullin the electronic charge toward itself an thereby increasing the electronic charg between the two. This is precisely wha happens in a covalent bond, i.e., buildu of electron charge density between th nuclei. The polarising power of the cation the polarisability of the anion and th extent of distortion (polarisation) of anio are the factors, which determine the pe Dipole moments of some molecules are cent covalent character of the ionic bond shown in Table 4.5. Just as all the covalent bonds have 4.4 The Valence Shell Electro some partial ionic character, the ionic Pair Repulsion (VSEPR) Theory bonds also have partial covalent character. As already explained, Lewis concept is unabl The partial covalent character of ionic to explain the shapes of molecules. Thi bonds was discussed by Fajans in terms of theory provides a simple procedure to predic the following rules: the shapes of covalent molecules. Sidgwic Table 4.5 Dipole Moments of Selected Molecules Dipole Type of Molecule Example Geometry Moment, µ(D) Molecule (AB) HF 1.78 linear HCl 1.07 linear HBr 0.79 linear Hl 0.38 linear H 0 linear 2 Molecule (AB ) H O 1.85 bent 2 2 H S 0.95 bent 2 CO 0 linear 2 Molecule (AB ) NH 1.47 trigonal-pyramidal 3 3 NF 0.23 trigonal-pyramidal 3 BF 0 trigonal-planar 3 | | | | | s e li g d e t p e , e n r . n e s t k | | --- | --- | --- | --- | --- | --- | | | Type of Molecule | Example | Dipole Moment, µ(D) | Geometry | | | | Molecule (AB) | HF HCl HBr Hl H 2 | 1.78 1.07 0.79 0.38 0 | linear linear linear linear linear | | | | Molecule (AB ) 2 | H O 2 H S 2 CO 2 | 1.85 0.95 0 | bent bent linear | | | | Molecule (AB ) 3 | NH 3 NF 3 BF 3 | 1.47 0.23 0 | trigonal-pyramidal trigonal-pyramidal trigonal-planar | | | | | | | | | | | Molecule (AB ) 4 | CH 4 CHCl 3 CCl 4 | 0 1.04 0 | | | **Table on page 24 (2×2)** | bond a atom is 121° er two in ethe m are | ngle is 117.6° . The formation ne is shown in | | --- | --- | | | | | | | **Table on page 29 (1×3)** | | | | | --- | --- | --- | | | | h r y |

• Matter consists of elements made of one type of atom.
• Atoms rarely exist independently except noble gases.
• Atoms combine to form molecules with distinct properties.
• Chemical bonding is the force holding atoms together.
• Atoms tend to achieve stable electronic configurations (octet rule).
• Kössel-Lewis approach explains bonding via electron transfer or sharing.
• 📌 Element: A substance made of only one type of atom.
• 📌 Molecule: A group of atoms bonded together with characteristic properties.
• 📌 Chemical Bonding: Force of attraction holding atoms in molecules or compounds.

Chemical Bonding

Chemical bonding is the force of attraction that holds atoms together in molecules or compounds. The Kössel-Lewis approach explains bonding in terms of atoms achieving stable electronic configurations, often resembling the nearest noble gas configuration. This is typically achieved by atoms either transferring electrons (ionic bonding) or sharing electrons (covalent bonding). The octet rule is central to this approach, stating that atoms tend to have eight electrons in their valence shell. However, this rule has limitations, especially for molecules involving elements with fewer or more than eight electrons around them. The chapter describes how atoms combine by either losing, gaining, or sharing electrons to complete their octets. The formation of ions and molecules is explained through this approach, which also introduces Lewis symbols to represent valence electrons. These symbols help visualize the bonding process by showing electron dots around atomic symbols. The section also highlights that the stability of molecules arises from the formation of chemical bonds, which lower the overall energy of the system.

• Chemical bonding holds atoms together in molecules or compounds.
• Kössel-Lewis approach focuses on atoms achieving stable electronic configurations.
• Atoms achieve stability by transferring or sharing electrons.
• Octet rule guides the formation of stable molecules.
• Lewis symbols represent valence electrons for bonding visualization.
• Bond formation lowers the energy and increases stability.
• 📌 Kössel-Lewis Approach: A method to explain bonding based on electron transfer or sharing.
• 📌 Lewis Symbol: Representation of valence electrons as dots around atomic symbols.
• 📌 Octet Rule: Atoms tend to have eight electrons in their valence shell.