Unit 1 - Chemical Bonding

Unit I -Chemical Bonding Valency of an element Valence is the number of bonds formed by an atom in a molecule. Valence electrons: The electrons in the outer energy level (i.e. valence shell) of an atom are the ones that take part in chemical bonding, and are referred as valence electrons. e.g. Sodium has electronic configuration as 2,8,1. and it forms only one chemical bond in its compounds hence its valence is one. While Chlorine has electronic configuration as 2,8,7. and it forms only one chemical bond in its compounds hence its valence is also one. Atoms with 8 electrons (or 2 electrons if only one orbit is there) in their valence shell are stable. Other atoms which do not have 8 electrons (or 2 electrons if only one orbit is there) in their valence shell, they try to attain the octate (or duplate if only one orbit is there) either by transferring some electrons or by sharing them and form chemical bonds. QB Chemical Bond: A chemical bond is defined as a force that acts between two or more atoms to hold them together so as to get a stable molecule (lower energy system). Based on mode of formation, chemical bonds can be classified as: I. II. III. IV. V. Ionic or Electrovalent Bond. Covalent Bond. Coordinate covalent Bond. Metallic Bond. van der Waals forces. (forces of attractions on molecular level) I. Ionic or Electrovalent Bond. QB Definition: The bond formed by transfer of an electron from one atom to another atom is called as Ionic or Electrovalent bond.e.g. NaCl, MgCl2, CaO. Example: Sodium Chloride is formed from an atom of sodium and one atom of chlorine, Na (2,8,1) has one valence electron while Cl(2,8,7) has seven valence electrons, Na transfers its one valence electron to Cl and both achieve noble gas configuration i.e. stable electron octate. Thus Na gives Na+ and Cl gives Cl- ion. QB Factors Governing formation of Ionic Bond: 1. Ionisation Energy : Ionisation energy of the metal atom, which loses electron should be low so that the formation of positively charged ion is easier. Lower the ionisation energy, greater is the tendency of a metal atom to change into cation and hence greater will be the ease of formation of ionic bond, thus alkali metals and alkaline earth metals forms ionic bonds easily. 2. Electron Affinity: The atom which accepts an electron and changes into an anion should have high electron affinity. Higher the electron affinity more is the energy released and stable will be the anion formed. Halogens and Gr VI A Elements tend to form Ionic bonds. ( less energy remains in the system, more stability) 3. Electronegativity difference of atoms involved: Electronegativity difference between atoms involved should be more than 2 only then they will form ionic bond. e.g Na has electronegativity 0.9 while Cl has 3.0, so the difference is 2.1 hence Na and Cl will form an Ionic compound. Characteristics of Ionic Compounds: 1. They are solids at room temperature. 2. They have high melting and boiling points. 3. They are hard and brittle 4. They are soluble in water. 5. They are conductors of electricity in molten and in solution state. 6. Reactions involving ionic compounds are fast. 7. Ionic bonds involve electrostatic lines of force between oppositely charged ions, it is not rigid and is non directional therefore ionic compounds are incapable of exhibiting isomerism. II. Covalent Bond. Defination: The attractive force between atoms, created by sharing of an electron pair is called as a covalent bond and the compouds containing covalent bond are called covalent compounds. Electronegativity difference of atoms involved should be less than 2 to form a covalent bond. Based on Electronegativity difference covalent bond can be categorised into two types. 1. Non polar covalent Bond A. If the electronegativity difference is 0 i.e. two atoms of same element are involved in bond formation and bond thus formed is called Non polar covalent bond. B. In such case bonded electrons are equally shared between two similar atoms. e.g. H2, Cl2, N2, F2, O2. QB 2. Polar covalent Bond A. If the electronegativity difference is between 0.1 to 2.0 i.e. two atoms of different elements are involved in bond formation and bond thus formed is called Polar covalent bond. e.g. HCl, NH3, H2O. H : Cl B. In such case bonded electrons are unequally shared between two different atoms and the electron pair is pulled towards more electronegative atom. C. Since electron pair is unequally shared, more electronegative atom aquires partial negative charge and less electronegative atom aquires partial positive charge thus molecule is referred as a dipole. D. Such polar compounds possess partial ionic character which can be calculated using the following formula. % ionic character = 16(XA- XB) +3.5 (XA – XB)2 ### 2. **Water (HO)** - **Electronegativity**: Oxygen (O) has an electronegativity of about 3.5, while hydrogen (H) is again about 2.1. The difference (1.4) indicates a significant polarity in the O-H bonds. - **Bond Formation**: In the O-H bonds, oxygen pulls the shared electrons closer to itself, resulting in a partial negative charge (-) on the oxygen atom and partial positive charges (+) on the hydrogen atoms. - **Molecular Geometry**: Water has a bent (or V-shaped) structure due to the two lone pairs on the oxygen atom. The asymmetry of this shape ensures that the dipoles from the O-H bonds do not cancel, leading to a net dipole moment. This makes HO a polar molecule as well. ### 1. **Ammonia (NH)** - **Electronegativity**: Nitrogen (N) has an electronegativity of about 3.0, while hydrogen (H) has an electronegativity of about 2.1. The difference (0.9) indicates that nitrogen is more electronegative than hydrogen. - **Bond Formation**: In the N-H bonds, nitrogen attracts the shared electrons more strongly than hydrogen does. This creates a partial negative charge (-) on the nitrogen atom and a partial positive charge (+) on each hydrogen atom. - **Molecular Geometry**: Ammonia has a trigonal pyramidal shape due to the presence of a lone pair on the nitrogen atom. This geometry means that the dipoles from the N-H bonds do not cancel out, resulting in a net dipole moment, making NH a polar molecule. E. Hydrogen Bonding: Hydrogen Bonding : When Hydrogen is covalently bonded to highly electronegative atom X, (O, F, N,), The shared electron pair is pulled so close to X, that a strong dipole results, Defination: The electrostatic attaraction between H atom covalently bonded to highly electronegative atom X, and a lone pair of electrons of X in another molecule is called hydrogen bonding. • Only O, N and F which have very high electronegativity and small atomic size are capable of forming hydrogen bonds. • Hydrogen bond is longer and much weaker than normal covalent bond. • Hydrogen bond results in long chain or cluster of large number of associated molecules. • Like a covalent bond hydrogen bond has a bonding direction. Types of Hydrogen Bonding: i. Intermolecular H Bonding: This type of bond is formed between two different molecule of same of diferent substance. e.g. Hydrogen bonding in HF, H2O, NH3. This type of H bond result in the formation of associated molecules, thus substances with intermolecular H bond have high melting points, boiling points, viscocity, surface tension etc. Ice (Solid HO): Hydrogen bonding between water molecules. Covalent bonds within each water molecule (O-H bonds). HF (Hydrogen Fluoride): Covalent bonds between H and F. Strong hydrogen bonding between HF molecules. Naphthalene (CH): Covalent bonds within the molecule. London dispersion forces between molecules (since it is nonpolar). Diamond (Carbon): Covalent bonds in a tetrahedral network (giant covalent structure) . HO (Water): Covalent bonds within each water molecule. Hydrogen bonding between water molecules . HS (Hydrogen Sulfide): Covalent bonds between H and S. Dipole-dipole interactions and London dispersion forces between molecules. Ethyne (CH): Covalent bonds within the molecule (C-H and CC bonds). London dispersion forces between molecules. HO (Hydronium Ion): Covalent bonds between H and O in the ion. Hydrogen bonding with surrounding water molecules in aqueous solution. HCl (Hydrogen Chloride): Polar covalent bonds between H and Cl. Dipole-dipole interactions between HCl molecules.