Chapter 11

1. Chapter 11 Liquids and solids

2. Almost all substances that are liquids are molecular, (held together by the covalent bonds within the molecule) • The physical properties of molecular liquids and solids is due to the intermolecular forces that hold them together

3. They are similar • compared to gases. • They are incompressible. • Their density doesn’t change with temperature. • These similarities are due • to the molecules being close together in solids and liquids • and far apart in gases • What holds them close together?

4. Intermolecular forces • Inside molecules (intramolecular) the atoms are bonded to each other. • Intermolecular refers to the forces between the molecules. • These are what hold the molecules together in the condensed states.

5. Intermolecular forces are much weaker than ionic or covalent bonds • Less energy is required to vaporize water, than to break the bonds between hydrogen and oxygen. • When a molecule changes state from solid-liquid-gas, the molecule itself stays together

6. Properties of liquids • Many are related to the intermolecular forces that hold them together • Boiling point- in order for a liquid to boil the molecules must overcome their attractive forces between them in order to separate. The stronger the intermolecular force, the higher the boiling point. (Melting point is based on the same relationship)

7. Ion dipole forces • Refer to a solution, not a molecule. • Force that exists between an ion and the partial charge on the end of a polar molecules. • Like NaCl dissolved in H2O • The strength of the force increases as the charge on the ion increases, or the strength of the dipole increases.

8. Types of Intermolecular Forces 3 Types for neutral molecules: Dipole-dipole London Dispersion Hydrogen All are based on electrostatic attractions All called Van der Waals forces (man who predicted the deviation of gases from the ideal gas laws due to attractive forces between them)

9. Dipole - Dipole • 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 • Only work when polar molecules are very close together

10. For molecules of approximately equal mass and size, the strengths of intermolecular attractions increase with increasing polarity • The higher the dipole moment, the higher the boiling point

11. London Dispersion Forces • Non - polar molecules also exert forces on each other. • Otherwise, no solids or liquids. • Electrons are not evenly distributed at every instant in time. • Have an instantaneous dipole. • Induces a dipole in the atom next to it. • Induced dipole- induced dipole interaction.

12. London Dispersion Forces • Weak, short lived. • Lasts longer at low temperature. • Eventually long enough to make liquids. • More electrons, more polarizable. • Bigger molecules, higher melting and boiling points. • Much, much weaker than other forces. • BP & MP of halogens and noble gases increase as you go down

13. 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.

14. H2O HF H2Te H2Se NH3 SbH3 H2S HI AsH3 HCl HBr PH3 SnH4 GeH4 SiH4 CH4 100 Boiling Points 0ºC -100 200

15. Water d+ d- d+

16. d- d+ d- d d- d+ H H H H H H H H H H H H Example

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

18. Viscosity • Which will pour very easily, syrup or lemonade? Both liquids so what accounts for the difference? • Viscosity- How much a liquid resists flowing. • The higher the viscosity, the slower it flows.

19. Related to how the molecules of the liquid move with respect to one another. • Viscosity decreases as temperature increases.

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

21. Water has a high surface tension because of it’s hydrogen bonds • Cohesive- the intermolecular forces that bind like molecules together (very strong for water) • Adhesive- the intermolecular forces that bind a molecule to another surface. • Mercury/ water

22. Capillary Action • Liquids spontaneously rise in a narrow tube. • Helps water move through plants • Occurs because of the attraction between the polar glass and the polar water molecules.

24. Beading • If a polar substance is placed on a non-polar surface. • There are cohesive, • But no adhesive forces. • When you wax car.

25. Relation to intermolecular forces? • The stronger the force, the bigger the effect. • Hydrogen bonding • Polar bonding • LDF

26. Phases • The phase of a substance is determined by three things. • The temperature. • The pressure. • The strength of intermolecular forces.

27. Phase changes • Directly related to the strength of the intermolecular forces> • Na has a boiling point of 883°C (strong metallic bonds) • While iodine readily sublimes at room temperature (LDF) • When your phase change is going to a less ordered state (more entropy) energy must be supplied

28. Energy Requirement • Fusion- to melt, the energy required to change from the solid to liquid state is called the heat of fusion (Hfus) • Vaporization- to turn from a liquid to a gas, the energy required to change from a liquid to a gas is called the heat of vaporization (Hvap). • The heat of vaporization is usually much larger than the heat of fusion.

29. Fusion, vaporization, and sublimation are endothermic • Freezing, condensation, and deposition are exothermic • Steam burns, sweating cools

30. Changes of state • The graph of temperature versus heat applied is called a heating curve. • The temperature a solid turns to a liquid is the melting point.

31. Heating Curve for Water Steam Water and Steam Water Water and Ice Ice

32. Slope is Heat Capacity Heat of Vaporization Heating Curve for Water Heat of Fusion

33. Calculating Energy Changes • Q=mcΔT • C= specific heat (given in J/g-K) • M=mass • ΔT= change in T, no need to convert • To melt To vaporize • molesHfus molesHvap

34. Calculate the enthalpy change of converting 1 mole of ice at -25°C to water vapor at 125°C. (The pressure is held constant at 1 atm). The specific heats of ice is 2.09 J/g-K, the specific heat of water is 4.18 J/g-K, and the specific heat of vapor is 1.84 J/g-K. The heat of fusion for water is 6.01 kJ/mol, and the heat of vaporization is 40.67 kJ/mol.

35. You Try! • Ethanol (C2H5OH) melts at -114°C, and boils at 78°C. The enthalpy of fusion of ethanol is 5.02 kJ/mol, and the enthalpy of vaporization 38.56 kJ/mol. The specific heats of solid and liquid ethanol are respectively .97J/g-K, and 2.3J/g-K. How much heat is required to convert 75 g of ethanol at -120°C to the vapor phase at 78°C?

36. Critical Temperature and Pressure • As temperature rises, gases become harder to liquefy because their molecules have high KE and are very far apart • Critical Temperature- the highest temperature a substance can exist as a liquid • Critical Pressure- The pressure required to liquefy the gas at the critical temperature. • The stronger the intermolecular forces, the easier a gas will liquefy.

37. Vapor Pressure • Vaporization - change from liquid to gas at boiling point. • Evaporation - change from liquid to gas below boiling point

38. Vaporization is an endothermic process - it requires heat. • Energy is required to overcome intermolecular forces pushing the molecules far apart into the gas phase.

39. Condensation • Change from gas to liquid. • Achieves a dynamic equilibrium with vaporization in a closed system. • What is a closed system? • A closed system means matter can’t go in or out.

40. Dynamic Equilibrium • The two opposing processes of condensation and vaporization are occurring at the same rate. • May look as if nothing is happening because there is no net change • The vapor pressure is the pressure exerted by its vapor when the liquid and vapor states of a substance are at dynamic equilibrium

41. In an open system • As the water evaporates, the molecules spread out and are not recaptured, no equilibrium is reached and the water continues to evaporate until there is nothing left. • Liquids that have high vapor pressure evaporate quickly and are called volatile (gasoline) • Vapor pressure increases with increasing T.

42. Boiling Point • Reached when the vapor pressure equals the external pressure. • Normal boiling point is the boiling point at 1 atm pressure. • Where will water boil faster, at the top of a hill or the bottom of a valley? • Where will pasta cook faster, at the top of a hill or the bottom of a valley?

43. Phase Diagrams. • A plot of temperature versus pressure for a closed system, with lines to indicate where there is a phase change. • Shows where equilibrium exists between different states of matter. • Can use it to predict the most stable phase of matter at a given T and P.

44. Solid Liquid C C B B A A Gas D D Pressure D 1 Atm D Temperature

45. Critical Point Solid Liquid Triple Point Gas Pressure Temperature

46. Solid Liquid Gas • This is the phase diagram for water. • The density of liquid water is higher than solid water. Pressure Temperature

47. This is the phase diagram for CO2 • The solid is more dense than the liquid • The solid sublimes at 1 atm. Pressure Liquid Solid 1 Atm Gas Temperature

48. Solids • Two major types. • Amorphous- those with much disorder in their structure. (Rubber and glass) • Crystalline- have a regular arrangement of components in their structure. (Diamond and Quartz)

49. Melting Points • The melting points of crystals are definite. • The melting points of amorphous solids vary.

50. 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.