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Weak Interactions Non-Covalent Interactions MATERIALS SCIENCE &ENGINEERING Anandh Subramaniam & Kantesh Balani Materials Science and Engineering (MSE)

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Presentation on theme: "Weak Interactions Non-Covalent Interactions MATERIALS SCIENCE &ENGINEERING Anandh Subramaniam & Kantesh Balani Materials Science and Engineering (MSE)"— Presentation transcript:

1 Weak Interactions Non-Covalent Interactions MATERIALS SCIENCE &ENGINEERING Anandh Subramaniam & Kantesh Balani Materials Science and Engineering (MSE) Indian Institute of Technology, Kanpur URL: home.iitk.ac.in/~anandh AN INTRODUCTORY E-BOOK Part of A Learner’s Guide

2  We will discuss here about intermolecular weak interactions  The strong ‘bonds’ are: Covalent, Ionic and Metallic Weak Interactions

3  Boiling point (being the temperature at which vapor pressure of substance equals the ambient pressure) is a better measure of non-covalent forces as compared to Melting point (which not only influenced by Attractive forces but also the crystal packing) How do we get a measure of non-covalent interactions? NaCl → 1413  C H 2 O → 100  C Ar →  186  C BrF → 20  C Boiling point Decreasing BP  EN = 4.0 F  2.8 Br = 1.2 Electronegativity difference Polar covalent F  1.91 Br +   EN = 3.5 O  2.1 H = 1.4 Polar covalent O  H+H+ H+  Schematics Ar

4 Inter-molecular Bonding Intra-molecular COVALENT Hydrogen bond Van der Waals Etc. , ,  … Dipole-dipole Dipole- Induced dipole Instantaneous dipole-induced dipole London Dispersion Relative strengths dispersion forces < dipole-dipole interactions < hydrogen bonds Ion-dipole Cation-Pi Pi-Pi The term ‘Van der Waals forces’ is sometimes used for a specific type (London Dispersion) rather than the class We will describe briefly a few of these (only) here Noncovalent Interactions: A Challenge for Experiment and Theory, Klaus Müller-Dethlefs and Pavel Hobza, Chem. Rev. 2000, 100, 143−167 Further reading

5  The covalent boding between a hydrogen atom and a strongly electronegative atom becomes ‘polar’- covalent  The ‘charged’ hydrogen ‘ion’ can be attracted to a electronegative atom, such as nitrogen, oxygen or fluorine  hydrogen bond should not be confused with a covalent bond to hydrogen.  Types of hydrogen bonds:  Intermolecular (between molecular)  Intramolecular (within a molecule)  E.g. of hydrogen bonding: water (responsible for the high boiling point of water compared to say H 2 S), DNA, partly responsible for the secondary, tertiary, and quaternary structures of proteins and nucleic acids, Polymers Hydrogen bond O  H+H+ H+  O  H+H+ H+  Hydrogen bond Schematics

6 Hydrogen bonded Ice crystal (hexagonal) [0001] Ice crystallizes in many polymorphic forms (12 crystal structures and 2 amorphous forms known)- we consider one form here Value Lattice parameter(s)a = 4.52 Å c = 7.37 Å Space Groupp6 3 /mmc  Packing fraction → ~0.34  Note the low packing fraction in spite of having the same space group as HCP crystals  c/a ratio → (very near ideal ratio of 1.633)

7 Dipole- Dipole interactions  In the covalent bonding between two atoms of very different electronegativity the bond becomes highly polar (introducing partial charges on the species)  This dipole can interact with other permanent dipoles  This interaction is stronger than dispersion forces F  Br +  F  Br +  Diplole-Dipole Interaction Schematics Van der Waals

8  Instantaneously generated dipole (due to asymmetry in electron charge distribution around the nucleus) on one atom leads to slight polarization of the atom (→ quantum induced instantaneous polarization)  This induces a dipole on the neighbouring atom (temporarily)  The force between these two dipoles is called the London dispersion forces  The force is very weak and is temporally varying  Can operate between non-polar molecules (H 2, Cl 2, CO 2 etc.)  The strength of the dispersion forces will increase with number of electrons in the molecule Instantaneous dipole-induced dipole London Dispersion Ar  Ar ++ Schematics

9 Ion-Dipole  Permanent dipole interacts with an ion.  This explains for example the solubility of NaCl in water.  The figure below shows the interaction of Na + and Cl  ions interacting with the permanent dipoles in a water molecule. Schematics


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