Chapter 5: Liquids and Solids

1. Chapter 5: Liquids and Solids The attractive and repulsive forces between molecules in liquids and solids are some of the most important forces in Chemistry • Why different liquids boil at different temperatures • Why some liquids flow and others don’t • Why diamonds are the hardest known substances • Why DNA twists into a double helix • In this chapter, we’ll study Intermolecular Forces: the forces that exist between molecules

2. Phases • Last chapter we studied gases • According to one of the tenets of the kinetic molecular theory • The molecules do not attract each other • In liquids and solids, because of the proximity of the molecules to each other, the intramolecular forces between molecules ARE significant • We’ll refer to liquids and solids as Condensed phases

3. Interion and Intermolecular Forces • Ion-Ion interactions are the strongest interactions • Example of an ion-ion interaction? • Let’s look at the various interactions given in the table

4. Ion-Dipole Interactions • Best example: Hydrated Ions • The polar character of the water molecule allows it to interact with cations or anions • We can describe the interaction energy: z = ion charge µ = Electric dipole moment r = distance

5. Hydrated vs Anhydrous Molecules • When salts crystallize out of water, they sometimes pull water molecules with them Na2CO3•10H2O CuSO4•5H2O • These are called hydrated salts • The Group I and II metal cations form much more hydrated salts than transition metal salts • These cations are smaller (usually) than transition metal cations Remember: Molar mass calculation includes the water molecules!

6. Ion-dipole interactions are strong for small, highly charged ions; one consequence is that small, highly charged cations are often hydrated in compounds

7. Dipole-Dipole interactions • Let’s look at the interactions between polar molecules The Dipole-Dipole interactions force some order in the solution

8. Dipole-Dipole interactions Dipole-Dipole interaction energy: µ1: Dipole moment of molecule 1 µ2: Dipole moment of Molecule 2 The fact that the distance is cubed means that the energy falls of much more rapidly than ion-ion interactions as the interacting species are separated

9. Which molecule has the higher boiling point: p-dichlorobenzene and o-dichlorobenzene Dipole moment for the molecules?

10. Which molecule has the higher boiling point: cis-dichloroethene or trans-dichloroethene?

11. London Forces • London Forces are attractive forces between non-polar molecules • These are 1 of the two weakest intermolecular forces • How do these interactions arise?

12. London Forces • The electron clouds are constantly shifting and sometimes the molecule gets a small dipole moment • Neighboring nonpolar molecules will have their electron clouds distorted and will form a dipole of opposite orientation • Then the process starts over (Dipole disappears and reforms) (1x10-16 sec to form and disappear)

13. 1: Polarizability of molecule 1 2: Polarizability of molecule 2 London Forces Very short range effects!! r6 !!!! • What determines Polarizability? • Large atomic radii • Low Zeff • High Polarizability = Large London Interactions

14. Let’s look at London Forces and Polarizability with respect to physical properties As we go down a group, the atomic radius increases and the melting and Boiling points increase (takes More energy)

15. London Forces and Molecular Shape • Because the London Force energy drops off VERY sharply as a function of distance, molecular shape is a major contributor to London Force energy Which has the higher boiling point?

16. Dipole - Induced Dipole • The presence of a molecule with a strong dipole moment can induce or create a dipole in a non-polar molecule • This depends on the strength of the dipole and the polarizability of the nonpolar molecule 1: Dipole moment of molecule 1 2: Polarizability of molecule 2

17. Hydrogen Bonding • A special type of dipole-dipole interaction • Hydrogen bonding only occurs between: N-H O-H and Lone pair e- on N, O, F F-H

18. Hydrogen Bonding • Hydrogen bonds are one of the most important interactions in biological systems • Hydrogen Bonds: • Hold proteins together • Allow DNA base pairs to match up • Allow structural polymers to interact Hydrogen bonds are the strongest type of intermolecular force

19. Liquids • The liquid phase is a unique phase midway between the solid phase and the gas phase • In the liquid phase, the molecules move past and around one another, but they don’t reach the separation distances achieved by molecules in the gas phase • There is not enough energy to overcome all of the intermolecular forces present in the system

20. Liquids and Short Range Order • In the liquid state, we can look at a molecule and its immediate neighbors and see the interactions occurring • However, we cannot say what the exact arrangement will be some distance away from where we are looking • This is an example of Short Range Order

21. Viscosity • You’ve heard of viscosity before, but probably only in commercials on television for motor oil • Viscosity is defined as the resistance of a liquid to flow • Viscosity is dependent on temperature to some extent • As we pump more thermal energy into the system, the molecules can break their intermolecular interactions and move about more easily • Viscosity Decreases as Temperature Increases

22. Viscosity: Examples • Oils: Carbon chains / London Forces • Water: Hydrogen Bonding • Sulfur: S8 ring breakage

23. Surface Tension • Why is the surface of a liquid smooth? • The intermolecular forces in the liquid • The tendency for liquids to form droplets is a direct result of surface tension • Water forms droplets • Nonpolar molecules like chloroform only have London Forces and do not form as large a droplet

24. “Wetting” • Water is “wet” because it can form hydrogen bonds with a surface and it does so with ease • Let’s think about this in terms of things you’ve seen before • Water on a waxy leaf: It beads up; It doesn’t “wet” the leaf • Water on Paper: The water “wets” the paper because the cellulose polymers in the paper have lots of -OH groups

25. Capillary Action • Capillary action is the rise of liquids up a narrow tube • Happens as a result of the sum of the adhesive forces and cohesive forces of the molecule and the glass • Adhesion: Forces/Interactions between a liquid and a surface • Cohesion: Intermolecular forces that keep molecules bound together

26. Solids • Solids are formed when temperatures are low enough to prevent molecules from moving around their neighbors • Molecules or atoms in solids vibrate/oscillate around a fixed position • All the atoms in a solid essentially move in unison

27. Types of Solids • Two major classifications: Crystalline and Amorphous • Crystalline Solids: All atoms, ions or molecules in the solid lie in an orderly array. Long range order exists within a crystalline solid • Amorphous Solids: The atoms, ions or molecules in an amorphous solid lie in a seemingly random arrangement. • Amorphous solids look like we took a snapshot of a liquid • Glass is an amorphous solid

28. Crystalline Solids • Molecular Solids: Assemblies of discrete molecules held together by intermolecular forces (eg: Quartz) • Network Solids: Atoms are covalently bonded to their neighbors throughout the extent of the solid (eg: Diamonds) • Metallic Solids: Metal cations held together by a sea of electrons • Ionic Solids: Built from the mutual attraction of anions and cations