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Chap 1-1 intra and intermolecular forces

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Chap 1-1 intra and intermolecular forces

  1. 1. Instrumental techniques Phar 6521 1
  2. 2.  Chromatography is a physical method of separation in which the components to be separated are distributed between two phases (KD/P = Distribution/partition constant)  one of which is stationary (stationary phase) while the other (the mobile phase) moves through it in a definite direction.  The chromatographic process occurs due to differences in the distribution constant of the individual sample components. 3 Chromatography
  3. 3. KD of Cpd A = [A]S / [A]M KD = Distribution constant of compound A [A]S = concentration of compound A in stationary phase [A]M = concentration of compound A in mobile phase For eg. TLC Chromatography  compounds distributes itself b/n a liquid mobile phase and a solid stationary phase  The rate of migration for a chemical compound is determined by how much of it distributes into the mobile and stationary phases Case 1. A compound that distributes itself 100% into the mobile phase  will migrate at the same rate of the mobile phase Case 2: On the other hand, a compound that distributes itself 100% in the stationary phase  will not migrate at all 4
  4. 4. In most molecular substances, there are two types of attractive forces: 1. Intramolecular and 2. Intermolecular forces 5  Chromatographic separation is achieved due to  Different compounds have different KD values  Distribution constant (KD) is affected by  the type of intermolecular forces that present in molecules
  5. 5.  force that hold atoms in a single molecule or a force of attraction within a molecule e.g. covalent bond, ionic bond H-Cl, Na + Cl - Intramolecular forces 6
  6. 6.  an attraction between two or more separate molecules.  are the result of attractions between positively and negatively charged regions of separate molecules.  They are not as strong as intramolecular force (chemical bonds). Intermolecular forces 7
  7. 7.  These intermolecular forces as a group are referred to as van der Waals forces. There are three types of intermolecular forces, 1. Dipole-dipole interactions 2. Hydrogen bond 3. London force/Dispersion force 8
  8. 8. A very approximate strength order would be: Bond type Relative strength Ionic bonds 1000 Hydrogen bonds 100 Dipole-dipole 10 London forces 1 9
  9. 9. Intermolecular Forces  They are, however, strong enough to control physical properties such as, solubility, boiling and melting points, vapor pressures, and viscosities. 10
  10. 10. Dipole-Dipole Interactions • Molecules that have permanent dipoles are attracted to each other. √ The positive end of one is attracted to the negative end of the other and vice-versa. √ These forces are only important when the molecules are close to each other. 11
  11. 11.  It occurs in polar compounds  These work in a similar manner to ionic interactions, but are weaker because only partial charges are involved.  An example of this can be seen in Acetone Dipole-Dipole Interactions 12 ----------
  12. 12. Dipole-Dipole Interactions The more polar the molecule, the higher is its boiling point. 13
  13. 13. Hydrogen Bonding Hydrogen bonding occurs when Hydrogen is bonded to N, O, or F are unusually strong. Hydrogen atom has a partial positive charge and can interact with another highly electronegative atom in an adjacent molecule (N, O, or F).  it is a special type of dipole-dipole force The result is a dipolar molecule  e.g H2O, NH3, HF 14
  14. 14. London/Dispersion Forces  While the electrons in the 1s orbital of helium would repel each other (and, therefore, tend to stay far away from each other), it does happen that they occasionally wind up on the same side of the atom.  It involve the attraction between temporarily induced dipoles 15
  15. 15. London/Dispersion Forces  At that instant, then, the helium atom is polar, with an excess of electrons on the left side and a shortage on the right side. 16
  16. 16. London/Dispersion Forces  Another helium nearby, then, would have a dipole induced in it, as the electrons on the left side of helium atom 2 repel the electrons in the cloud on helium atom 1. 17
  17. 17. London/Dispersion Forces  London dispersion forces, or dispersion forces, are attractions between an instantaneous /temporary dipole and an induced dipole. 18
  18. 18. London/Dispersion Forces These forces are present in all molecules, whether they are polar or non-polar. The tendency of an electron cloud to distort in this way is called polarizability. 19
  19. 19.  This polarization can be induced either by  A polar molecule or  A non-polar molecule (the repulsion of negatively charged electron clouds in non-polar molecules) London/Dispersion Forces 20
  20. 20. 21
  21. 21. Factors Affecting London Forces The strength of dispersion forces tends to increase with increased molecular weight. Larger atoms have larger electron clouds, which are easier to polarize. 22
  22. 22. Factors Affecting London Forces The shape of the molecule affects the strength of dispersion forces: long, skinny molecules (like n-pentane tend to have stronger dispersion forces than short, fat ones (like neopentane). This is due to the increased surface area in n-pentane. 23
  23. 23. Which Have a Greater Effect: Dipole-Dipole Interactions or Dispersion Forces? • If two molecules are of comparable size and shape, dipole-dipole interactions will likely be the dominating force. • If one molecule is much larger than another, dispersion forces will likely determine its physical properties. 24
  24. 24. Ion-Dipole Interactions 25 • A fourth type of force, ion-dipole interactions are an important force in solutions of ions. • The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents.
  25. 25. Anthocyanins (also anthocyans) are  belong to a parent class of molecules called flavonoids  cationic organic compound  well water-soluble pigments due to ion-dipole interaction
  26. 26. Quaternary ammonium cation are salts of quaternary ammonium cations. Soluble in water
  27. 27.  Polar Molecule A Molecule with a Positive and Negative Side Dipole Moment • A Measure of Molecular Polarity • A Non-polar Molecule will have a Zero Dipole Moment 28 Molecular Polarity
  28. 28. Why is Polarity Important?  Many Properties Depend on Polarity Melting and Boiling Point Surface Tension, Viscosity Reactivity Solubility (e.g., will it dissolve in water) 29
  29. 29. Requirements  A Polar Molecule Requires Polar Bonds A Molecular Shape that Separates the by the partial Positive from the Negative Side 30
  30. 30.  Unequal sharing of e- in a bond called Polar Covalent Bond or Polar Bond  Partial Charge indicated by delta +/- Polar Covalent Bonds -+ 31
  31. 31. Polar Covalent Bonds  NON-polar covalent bonds  Bonds between identical atoms such as H- H, F-F involve equal sharing of e-  Polar covalent bonds  Bonds between different atoms involve unequal sharing of e-  Polar Covalent Bonds have a Partial Charge Separation 32
  32. 32. Electronegativity  Electronegativity is the ability of an atom to attract electrons in a bond.  Nonmetals bonded to N, O, F usually have polar bonds 33
  33. 33. Electronegativity Index of Some Elements 34
  34. 34. Electronegativity  Used to Determine Bond Polarity   EN < 0.45 Non-polar Bond  1.75 >  EN > 0.45 Polar Bond   EN > 1.75 Ionic Bond Atoms 35 Polar or Non-polar Bond? O-H  EN = 1.4 Polar Bond C-H  EN = 0.4 Non-polar Bond C-O  EN = 1.0 Polar Bond H-Cl  EN = 0.8 Polar Bond
  35. 35. Polar Molecules  The molecule is usually polar  If all atoms attached to central atom are not the same  Or if central atom has 1 or more lone pairs of electrons 36
  36. 36. Examples  Polar compounds  HCN, H2O, CHCl3, CH2Cl2  Non-polar  CO2, CCl4 37
  37. 37. In Conclusion  Polarity Determined from Polar Bonds ( EN > 0.45) Molecular Shape with + & - sides N and O in Molecules often lead to Polar Molecules or Regions 38
  38. 38. In Conclusion  Polarity will be important  In Determining Intermolecular Forces Vapor Pressure, Boiling and Melting Points Surface Tension, Viscosity  Solubility  Reactivity (Organic Chemistry) 39
  39. 39. Intermolecular Forces Affect Many Physical Properties The strength of the attractions between particles can greatly affect the properties of a substance or solution.
  40. 40. Solubility  Defines as the amount of a solute that will dissolve in a specific solvent at given condition 41 Degree of solubility (types of saturation) Saturated solution: A solution with solute that dissolves until it is unable to dissolve anymore, leaving the undissolved substances at the bottom. Unsaturated solution: A solution (with less solute than the saturated solution) that completely dissolves, leaving no remaining substances. Supersaturated solution: solution (with more solute than the saturated solution) that contains more undissolved solute than the saturated solution because of its tendency to crystallize and precipitate
  41. 41. Factors that affects solubility  The nature of solute and solvent  Temperature  Pressure (only applicable to gases) 42
  42. 42. The nature of solute and solvent  When two substances are similar they can dissolve in each other Polar solutes dissolve in polar solvent Non-polar solutes tend to dissolve in non-polar solvent  “Like dissolve like”  two liquids dissolve in each other b/c the molecules are alike in polarity 43 Note: solvents are grouped either polar or non-polar solvent Polar Solvent: a liquid made up of polar molecules Non-polar Solvent: a liquid made up of non-polar molecules
  43. 43. The nature of solute and solvent  Ionic compounds are made up of charged ions similar to polar compounds e.g. NaCl  Ionic compounds are more soluble in polar solvent than in a non-polar solvent  ion-dipole interaction 44
  44. 44. The nature of solute and solvent 45
  45. 45. Temperature  Solubility of solids in liquids  The solubility of a solid increases as temp increases Solubility of gases in liquid are affected by temperature  Opposite to the solubility of solids in liquids As the temperature increases, the solubility of a GAS in a liquid decreases WHY ?  As the temperature increases, the kinetic energy of the solute gas increases and the gas can escape 46
  46. 46. Pressure  when the pressure is increased over the solvent, the solubility of gas is increased  WHY ?  Pressure increases as gas molecules strike the surface to enter the solution increased 47
  47. 47. Factors of Dissolving  Rate of which a solid solute dissolves in a solution depends on three factors  Surface area: speed up the solubility by increasing surface area  Stirring: increases contact b/n solvent and solute  Temperature: kinetic energy increase 48

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