Chemical Bonding PDF

# Chemical Bonding ## Chemical Band - **Strong Bonds (~200 KJ)** - Covalent Bond - Ionic Bond - Co-ordinate Bond - Metallic bond - **Weak Bonds (2 to 40 KJ)** - Vanderwaals Force (2 to 10 KJ) - Hydrogen Bond (10 to 40 KJ) ### Description of Bonds - **Covalent Bond:** Mutu...

# Chemical Bonding ## Chemical Band - **Strong Bonds (~200 KJ)** - Covalent Bond - Ionic Bond - Co-ordinate Bond - Metallic bond - **Weak Bonds (2 to 40 KJ)** - Vanderwaals Force (2 to 10 KJ) - Hydrogen Bond (10 to 40 KJ) ### Description of Bonds - **Covalent Bond:** Mutual sharing of electrons. - **Ionic Bond:** Complete transfer of electrons. - **Co-ordinate Bond:** Shared electron pair is contributed by only one atom. - **Metallic Bond:** Force between Kernel and Free electrons. ## Lewis Dot Structure 1. Identify the central atom and surrounding atom. 2. Assign negative charge to more electronegative atom and positive charge. 3. Oxygen atom forms two bonds, but oxygen forms one bond. 4. Try to complete octet and duplet of all atoms. 5. But if central atom contains vacant d-orbital such as P, S, Cl, etc, then the octet can expand. ## Examples of Lewis Dot Structures 1. **C104** ``` :O: || O = Cl = O || :O: ``` 2. **CO32-** ``` :O: || O-C-O || :O: ``` 3. **NO3-** ``` :O: || O-N=O || :O: ``` ## Other Examples 1. **HND2** ``` H-O-N=O | H ``` 2. **Other Examples** - SO42- - PO43- ## Formal Charge *Formal Charge = Valence e - lone pair e - no. of surrounding bonds.* ## Bond order From Resonance *B.O = Total no. of Bonds / Total no. of positions.* ## Average Formal charge on a surrounding atom *Average Formal Charge = Net charge on species / no. of surrounding atoms* ## Valence Bond Theory (VBT) 1. A covalent Bond is a result of overlap of atomic orbitals of two atoms. 2. The atomic orbitals which overlap must contain unpaired electrons with opposite spins. Also, they must possess either same or nearly same energies. 3. Strength of covalent bond & extent of overlap. ## Types of Overlap 1. **S-S overlap:** Always form σ bond. 2. **p-p overlap:** Can Form σ and π bond 3. **s-p overlap:** Always form σ bond ## Strength Order *1s-1s > 1s-2p > 2p-2p > 2s-2p > 2s-2s > 3P-3P > 3s-3P > 3s-3s* ## Concept of Hybridisation (modified VBT) The mixing and recasting of atomic orbitals of same atom to form new orbitals called hybrid orbitals which are identical in all respect. ## Covalent Bond is a result of any of the Following overlap 1. Hybrid orbital of one atom and atomic orbital of another atom. (ex: CH4 Formation) 2. Hybrid orbital of one atom and hybrid orbital of another atom. (ex: Graphite) ## Types of hybridisation 1. **sp - (2)** - Linear 2. **sp2 - (3)** - Trigonal planar 3. **sp3 - (4)** - Tetrahedral 4. **sp3d - (5)** - Trigonal Bipyramidal. 5. **sp3d2 - (6)** - Octahedral. 6. **sp3d3 - (7)** - Pentagonal Bipyramidal. 7. **dsp2 - (4)** - Square planar 8. **d3s - (4)** - Tetrahedral. 9. **dsp3 - (5)** - Trigonal Bipyramidal. 10. **d2sp3 - (6)** - Octahedral. 11. **d3sp3 - (7)** - Pentagonal Bipyramidal. 12. **sp2d - (4)** - Square planar. ## Method of Finding hybridisation - *X* = Indicate valence e⊕ of central atom. - If surrounding atom is monovalent (H, F, Cl, Br, I, etc) then allot 1 ‘X’ form 1 atom. - If surrounding atom is divalent (O, S) then allot 2 ‘X’ Form 1 atom. ## Valence Shell Electron Pair Repulsion (VSEPR Theory) 1. If central atom is surrounded by only bond pairs and no lone pairs then geometry and shape both are same. ex. CH4. 2. If central atom is surrounded by both bond pairs and lone pairs then geometry and shape both are different. 3. The order of repulsion of various types of e⊕ pairs is lp-lp > lp-bp > bp-bp. 4. Lone pair of central atom reduces bond angle from expected value. Decrease in bond angle from expected value α no. of lone pairs on central atom. ## Code Number, Hybridization, Geometry and Shape Table | Code Number | Hybridization | Geometry | Shape | |-------|--------------|-----------|------------| | 220 | sp | Linear | Linear (O2, BeCl2) | | 330 | sp2 | Trigonal planar | Trigonal planar (BF3) | | 321 | sp2 | | V-shape (SnCl3) | | 440 | sp3 | Tetrahedral | Tetrahedral (CH4) | | 431 | sp3 | | Pyramidal (NH3) | | 422 | sp3 | | V-shape (H2O) | | 550 | sp3d | TBP | TBP (PCl5) | | 541 | sp3d | | See-Saw (SF4) | | 532 | sp3d | | T-shape (ClF3) | | 523 | sp3d | | Linear (XeF2) | | 660 | sp3d2 | Octahedral | Octahedral (SF6) | | 651 | sp3d2 | | Square pyramidal (BrF5) | | 642 | sp3d2 | | Square planar (XeF4) | | 770 | sp3d3 | PBP | PBP (IF7) | | 761 | sp3d3 | | Distorted octahedral | | 752 | sp3d3 | | Pentagonal planar | ## Bond parameters (measurable properties of covalent bond) 1. **Bond angle** - angle between two adjacent covalent bonds. - Hybr expected B.A - sp: 180 - sp2: 120 - sp3: 109°28' - B.A & %S character (H is diff) - sp: 50%S - sp2: 33.33% S - sp3: 25% S 2. **No. of lone pairs** - **Same** - Hybridization, surrounding atom - **Different** - central atom, no. of lone pairs then - B.A α 1/ no. of lone pairs - ex: CH4 > NH3 > H2O 3. **Same** - Hybridization, surrounding atom, lone pair. - **Different** - central atom, - B.A α EN of central atom < NH3 > PH3 > ASH3> S6H3 - H2O H2S > H2Se > H2Te 4. **Same** - Hybridization, central atom, lone pair - **Different** - Surrounding atom - B.A α EN of surr. atom. - F20 < C120 < BY20 ## Bond order - Bond order: The number of bonds between two atoms, calculated by using any of the following: - **MOT** - Total e ≤ 20, species is diatomic. - **Resonance** - B.D = total no. of Bonds / total no. of positions. - **Use structure**. ## Bond energy - Bond energy: Energy required to break 1 mole of particular bonds. - Bond energy α Bond multiplicity - Bond energy α %s character. - Bond energy x (1/ atomic size) - Bond energy x Bond polarity - Bond energy α no. of lone pairs (use this factor when there is single bond between 2 identical atoms. ## Big exception - Big exception: in halogens - Expected order : F-F > CI-CI > Br-Br > I-I - Actual order: CI-CI > Br-BY > F-F > I-I ## Bond length - Bond length: The distance between the centres of nuclei of two closely bonded atoms. - BL α Bond multiplicity - BL α Atomic size - BL α %s character - BL α 1/ Bond order - ex: O2, O3, H2O2 - B.O (it bao ta yd talant) Bond length: Hin (d) - CO - CO2 - CO32- ## Dipole moment - Dipole moment: Measure of polarity of bond. - u = q x r - Units: esu.cm, C.m, Debye. - 1D = 10-18 esu - 1D = 3.33 X 10-30 cm. - MR = √ M² + 시½² + 2M, 42 COSO - % ionic character = u observed x 100 / u theory - Polar molecule = MR ≠ 0 (permanant dipole moment) - Non-polat molecule = MR = 0 ## Molecular orbital theory (Applicable only for diatomic species having totale ≤ 20) 1. During molecule or ion formation through covalent bonding, Atomic orbitals combine together to form new orbitals called molecular orbitals. 2. The number of molecular orbitals formed, are exactly equal to number of Atomic orbitals combined. 3. Two types of Molecular orbitals formed: - **Bonding MO.** (denoted by σ and π) - **Anti-Bonding MO.** (denoted by σ* and π*) 4. AD’S combined: | BMO’S combined: | ABMO'S formed: | |----|----|----| | ns + ns | σns | σ*ns | | npz + npz | σnpz | σ*npz | | пpx + пpx | πnpæ | π*npæ | | пpy + пpy | πnpy | π*npy | ## Filling of electrons in MO’S 1. **Configuration I:** when total e≤ 14. - ex: N2, N2+, B2, C2 etc (2s-12p mixing occurs) - σ1s < σ*1s < σ2s < σ*2s < σ2pz < π2px = π2py < σ2pz < π*2px = π*2py < σ*2pz. 2. **Configuration II:** when total ee = 15 to 20 ex: O2, O2-, F2, Nez etc (2s-2P mixing do not occurs) ## Applications of MOT 1. **Bond order** - B.O = 1 (No-Na) - Total e= | B.O= | - 1 | 0.5 | - 2 | 1.0 | - 3 | 0.5 | - 4 | 0 | - 5 | 1.0 | - 6 | 0.5 | - 7 | 1.0 | - 8 | 0 | - exception: Co+ = 3.5 2. **Bond length** - Bond length α 1/ Bond order 3. **Bond energy** - Bond energy & Bond order. - Thermal stability (thermal stability increases as bond order increases) 4. **Prediction of magnetic nature:** - Total e⊕ are odd or 10 or 16 → Paramagnetic - Total e⊕ are even → Diamagnetic ## Ionic Bond ( The Bond which is formed by complete transfer of e⊕ from one atom to another atom). - ΔΕΝ ≥ 1-6 - Low IE of element forming cation. - High electron: gain enthalpy with -ve sign of element forming anion. - High Lattice enthalpy. - Overall decrease in energy of system. ## Imp. note - **The Bond Formed between metal and non-metal is generally ionic.** - **The Bond Formed between two complex ions or in between simple ion and complex ion is also ionic.** - **There are some substance which contains ionic, covalent as well as co-ordinate bond.** ex: NH4Cl, [Ni(H2O)6]SO4. ## Electrovalency - Electrovalency: no. of electron lost or gained by one atom ## Lattice energy or lattice enthalpy (∆LH) Amount of energy released when 1 mole of ionic compound is formed from constituent gaseous cation and gaseous anion. - ∆LH = -ve. - or Amount of energy required to break 1 mole of ionic compound into its constituent gaseous cation and gaseous anion. - ∆LH = +ve. ## V. Imp: - Stability of ionic compound α ∆LH. - Solubility of ionic compound in H₂O α (1/ ∆LH). - ∆LH α 9c.ga - Yc+Ya ## Covalent nature in ionic bond: (concept of polarisation) - For cation → polarising power. - For anion → polarisibility - Covalent character in ionic substance α extent of polarisation ## Fajan's rules: - **Polarisation α size of anion** - Small cation <-> Large anion - **Polarisation α size of cation** - High charges - **Polarisation α 9c and ga.** - **Polarising power is exceptionally more if cation possess pseudo inert gas configuration (nsanpnd 19).** - ex: Nax <cu'x, KX < Agx ## V.Imp - Melting point, thermal stability, ionic nature, solubility in H2O α polarisation