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Intermolecular Forces

Intermolecular Forces. SCH4U1 October 3 2011 Mr. Dvorsky. Working with solids. At the macroscopic level, a solid is a substance that has both a definite shape and volume.

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Intermolecular Forces

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  1. Intermolecular Forces SCH4U1 October 3 2011 Mr. Dvorsky

  2. Working with solids • At the macroscopic level, a solid is a substance that has both a definite shape and volume. • At the microscopic level, solids are structures in which the particles are very close to each other. The intermolecular forces have overcome the kinetic energy of the particles, which exhibit no real motion, just vibration.

  3. Solids may crystalline • These solids display a very regular ordering of the particles in a 3D structure called a crystal lattice. In this lattice there are repeating units. • Includes: Atomic, Molecular, Network, Ionic, and Metallic.

  4. …or solids may be amorphous • These solids lack extensive ordering of the particles so the structure has a lack of regularity. They have indistinct shapes. • Includes: glass, polymers, charcoal.

  5. Intra versus Inter • Intramolecular forces are between atoms in a molecule (e.g. covalent, polar covalent). –these tend to be very strong. • Intermolecular forces are between molecules (e.g. London dispersions, dipole-dipole forces, hydrogen bonding). – tend to be very weak.

  6. Let us start with atomic solids • At very low temperatures, the noble gases form atomic solids and liquids • They already have the electron configuration other elements achieve during chemical bonding, therefore very stable and we would expect there to be little attraction between two atoms of noble gases. • Proof of this: noble gases can be liquefied or solidified only at extremely low temperatures.

  7. That being said, the fact that noble gases can form solids and liquids means there must be some forces of attraction between their atoms. -proof of the existence of intermolecular forces.

  8. In a noble gas like He, an instantaneous charge imbalance can occur. –that means at a given instant, both He electrons may be found in the same side of the atom • That side becomes slightly negative and the opposite side becomes slightly positive.

  9. London dispersion force • this is a temporary fluctuating dipole • the atom can then induce a dipole in a neighbouring atom and they can they attract each other • called London Force or London Dispersion Force.

  10. London dispersion force • It is a very weak force (1-10 kJ/bol)

  11. The strength of the London force depends upon size of the atom. • The larger the atom, the greater the number of electrons, the greater chance there is for a temporary imbalance of electrons and therefore the greater strength of force between that atom and neighbouring ones.

  12. With non-polar molecules, the force of attraction between them are these London forces that we discussed. • Therefore the only force of attraction between molecules is one which results from a temporary charge imbalance of electrons in the atoms which make up the molecules

  13. When comparing non-polar molecules with the same number of atoms, the size of the forces of attraction depends on the size of the atoms. • As the size of atom increases, and thus the size of molecule increases, the number of electrons increases too. Therefore more chance of temporary charge imbalance arising making the forces of attraction stronger.

  14. When comparing non-polar molecules in which the atoms are the same size but there are a different number of them, the forces of attraction are greater between molecules with the greatest number of atoms. • More atoms = more locations for LD forces to occur in adjacent molecules

  15. When comparing non-polar molecules in which the same number and type of atoms are present, the forces of attraction are greater between molecules with the shape that allows for more places where forces of attraction can arise between adjacent molecules

  16. Dipole-Dipole • The dipole-dipole force is the attraction between dipoles. • Recall a polar molecule has dipoles (oppositely charged ends). • Opposite charged ends of polar molecules attract

  17. In polar molecules, the forces of attraction are London Forces AND the dipole-dipole force. • Polar molecules such as HCl have a permanent dipole. • The slightly negative end of one molecule can attract the slightly positive end of a neighbouring molecule.

  18. Intermolecular Forces • The more polar the molecule (the larger the electronegativity difference), the stronger the dipole – dipole force and therefore the higher the boiling point. • e.g. ClF has electronegativity difference of 1.0, and has 26 electrons, boiling point: -101 C. BrF has EN difference of 1.2 and has 44 electrons, boiling point -20 C. -BrF has a greater dipole moment and a greater # of electrons thus dipole-dipole forces are greater, London dispersion forces are greater, and thus a higher boiling point.

  19. Account for the difference in BP between ClF and CH3Cl Account for the difference in BP between C2H6and CH3F Account for the difference in BP between ClF and CH3H8 Note: Sometimes you will run into a term called Van der Waals Forces, it is an all encompassing term that includes both London dispersion and dipole-dipole.

  20. Hydrogen Bonding • When hydrogen is covalently bonded to a very small highly electronegative atom such as fluorine, oxygen, or nitrogen, unusually strong dipole-dipole attractions are observed • Why? Two reasons.

  21. Because of the large electronegativity difference, the F-H, O-H, N-H bonds are very polar. Thus, the ends of these dipoles carry a substantial fraction of one charge • Because of the small size of the atoms involved, the charge on the end of a dipole is highly concentrated. -this makes it particularly effective at attracting the end of opposite charge on a neighbouring dipole. -these two factors combine to produce attractions called hydrogen bonds. -hydrogen bonds are 1/10th as strong as covalent bonds but 10x stronger than normal van der Waals forces.

  22. According to the trend, HF should have a BP less than -83.7. Why is its BP the highest?

  23. -If a substance has molecules with hydrogen bonds between them, it tends to have a higher melting point & boiling point. -hydrogen bonding also causes substances to form unusual structures in the solid state. When water freeze it forms an open expanded structure, therefore water expands when it freezes and ice has a lower density than liquid water. Thus ice floats. – for most substances solid state = more dense

  24. Hydrogen Bonding • Why is this important? -if ice was more dense than water, ice forming at the top of a body of water would freeze and fall to the bottom, thus it would freeze solid. -aquatic life could not survive -therefore we would have not evolved

  25. Hydrogen bonding in living systems • the structure of proteins is controlled by hydrogen bonding • Hydrogen bonding is an important factor in determining the shape of DNA • Plant fibres are more rigid than animal tissue because of more hydrogen bonding • We eat and wear things that are hydrogen bonded

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