Liquids and Solids

1. Liquids and Solids Chapter 10 – Day 1 Notes

2. L&S are similar to each other • Different than gases. • They are incompressible. • Their density doesn’t change much with temperature. • These similarities are due • to the molecules staying close together in solids and liquids • and far apart in gases • What holds them close together?

3. Intermolecular forces • Inside molecules (intramolecular) the atoms are bonded to each other. • Intermolecular refers to the forces between the molecules. • Holds the molecules together in the condensed states. • We already talked about these!!

4. Relative Magnitudes of Forces The types of bonding forces vary in their strength as measured by average bond energy. Strongest Weakest Covalent bonds (400 kcal/mol) Hydrogen bonding (12-16 kcal/mol) Dipole-dipole interactions (2-0.5 kcal/mol) Londonforces (less than 1 kcal/mol)

5. Dipole - Dipole • Remember where the polar definition came from? • Molecules line up in the presence of a electric field. The opposite ends of the dipole can attract each other so the molecules stay close together. • 1% as strong as covalent bonds • Weaker with greater distance. • Small role in gases.

6. + + - - - + + - + - - + + - - + - + + - Dipole - Dipole

7. Hydrogen Bonding • Especially strong dipole-dipole forces when H is attached to F, O, or N • These three because- • They have high electronegativity. • They are small enough to get close. • Effects boiling point.

8. Hydrogen Bonding Bonding between hydrogen and more electronegative neighboring atoms such as oxygen and nitrogen Hydrogen bonding between ammonia and water

9. Water d+ d- d+ Each water molecule can make up to four H-bonds

10. London Dispersion Forces The temporary separations of charge that lead to the London force attractions are what attract one nonpolar molecule to its neighbors. London forces increase with the size of the molecules. Fritz London 1900-1954

11. London Dispersion Forces

12. London Forces in Hydrocarbons

13. Let’s recap how to identify the IMF! • Hydrogen Bonding: Polar molecules containing H & N, O, or F. • Dipole – Dipole: All other Polar molecules. • London Dispersion Forces – Non-polar molecules

14. Boiling point as a measure of intermolecular attractive forces

15. Liquids • Many of the properties due to internal attraction of atoms. • Surface tension • Capillary action • Beading • Viscosity • Stronger intermolecular forces cause each of these to increase.

16. Surface tension • Molecules at the top are only pulled inside. • Molecules in the middle are attracted in all directions. • Minimizes surface area.

17. Capillary Action • Liquids spontaneously rise in a narrow tube. • Intermolecular forces are cohesive, connecting like things. • Adhesive forces connect to something else. • Glass is polar. • It attracts water molecules through cohesive action.

19. Beading • If a polar substance is placed on a non-polar surface. • There are cohesive, But no adhesive forces.

20. Viscosity • How much a liquid resists flowing. • Large forces, more viscous. • Large molecules can get tangled up. • Cyclohexane has a lower viscosity than hexane. • Because it is a circle- more compact.

21. Solids • Two major types. • Amorphous- those with much disorder in their structure. • Crystalline- have a regular arrangement of components in their structure.

22. Types of Solids • Amorphous solids: considerable disorder in their structures (glass).

23. Types of Solids • Crystalline Solids: highly regular arrangement of their components

24. Crystals • Lattice- a three dimensional grid that describes the locations of the pieces in a crystalline solid. • Unit Cell-The smallest repeating unit in of the lattice. • Three common types.

25. Cubic

26. Body-Centered Cubic

27. Face-Centered Cubic

28. Representation of Components in a Crystalline Solid • Lattice: A 3-dimensional system of points designating the centers of components (atoms, ions, or molecules) that make up the substance. • X-ray beams are shot through the solid to determine the arrangement of the atoms. • Bragg’s Law is the equation used to from the angles of the incident and reflective rays.

29. Bragg’s Law xy + yz = n and xy + yz = 2d sin n = 2d sin 

30. n = 2d sin  •  = wavelength • n = integer (1, 2, 3, …) • d = distance between atoms •  = angle of incidence and reflection

31. Example: • X rays of wavelength 1.54Å were used to analyze an aluminum crystal. A reflection was produced at θ = 19.3 degrees. Assuming n = 1, calculate the distance (in pm) between the planes of atoms producing this reflection. ( 1 Ångstrom = 100 pm)

32. Network Atomic Solids Some covalently bonded substances DO NOT form discrete molecules. Diamond, a network of covalently bonded carbon atoms Graphite, a network of covalently bonded carbon atoms

33. Molecular Solids Strong covalent forces within molecules Weak covalent forces between molecules Phosphorus, P4 Sulfur, S8