Unit 13 – Liquids and Solids

1. Unit 13 – Liquids and Solids 13.1 Water and Its Phase Changes 13.2 Energy Requirements for the Changes of State 13.3 Intermolecular Forces 13.4 Evaporation and Vapor Pressure 13.5 The Solid State: Types of Solids 13.6 Bonding in Solids Text Pages 398 - 416

2. Unit 13 – Liquids and Solids Upon completion of this unit, you should be able to do the following: • Use heat of fusion and heat of vaporization in energy calculations. • Describe dispersion interaction forces, dipole-dipole attraction, hydrogen bonding and London dispersion forces. Give examples of substances that exhibit these forces. • Explain the relationship among vaporization, condensation and vapor pressure • Explain the differences in physical properties, such as boiling points and rates of evaporation, in terms of strength of the forces of attraction. • List the characteristic properties of various types of crystalline structures

3. Water and Its Phase Changes • Pure water is a colorless, tasteless substance that at 1 atmosphere of pressure freezes to form a solid at 0 oC and vaporizes completely to form a gas at 100 oC. • The liquid range of water is 0 oC to 100 oC. • When liquid water is heated, the temperature of the liquid rises until it reaches 100 oC. It then stays at 100 oC until all the water has changed to vapor. Then the temperature will begin to rise again. • A graphical representation is shown in the next slide. This is called the heating/cooling curve for water. • Going from left to right, energy is added (heating). • Going from right to left, energy is being removed (cooling).

4. Water and Its Phase Changes

5. Water and Its Phase Changes • Water expands when it freezes. One gram of ice at 0 oC has a greater volume than one gram of water at 0 oC. • Water in a confined space can break its container when it freezes and expands. • The expansion of water when it freezes also explains why ice cubes float. • The density of water is 1.00 gram/ml • The density of ice is 0.917 gram/ml

6. Energy Requirements for Changes of State • Changes of state are physical, not chemical, changes. • No chemical bonds are broken in these processes. • The bonding forces that hold a molecule together are called intramolecular (within the molecule) forces. • The forces that occur among molecules that cause them to form a solid or a liquid are called intermolecular (between the molecules) forces. • It takes energy to melt ice and vaporize water because intermolecular forces must be overcome. • It would take much more energy to overcome the covalent bonds and decompose the water molecules into their component atoms.

7. Energy Requirements for Changes of State

8. Energy Requirements for Changes of State • The energy required to melt 1 mol of a substance is called the molar heat of fusion. • For ice, the molar heat of fusion is 6.02 kJ/mol. • The energy required to change 1 mol of liquid to its vapor is called the molar heat of vaporization. • For water, the molar heat of vaporization is 40.6 kJ/mol at 100 oC.

9. Energy Requirements for Changes of State Solid → Liquid Calculate the energy required to melt 8.5 g of ice at 0 oC The molar heat of fusion of ice is 6.02 kJ/mol. 2.8kJ

10. Energy Requirements for Changes of State Liquid → Gas Calculate the energy required to heat 25.0 g of liquid water from 25 oC to 100 oC and change it to steam at 100 oC. The specific heat capacity of water is 4.18 J/g-oC and the molar heat of vaporization of water is 40.6 kJ/mol. 7.8 + 57 = 65 kJ

11. Intermolecular Forces • Water is a polar molecule. It has a dipole moment. • When molecules with dipole moments are put together, they orient themselves to take advantage of their charge distributions. • Molecules with dipoles moments can attract each other by lining up so that the positive and negative ends are close to each other. This is called dipole-dipole attraction. • In a liquid, dipoles find the best compromise between attraction and repulsion.

12. Intermolecular Forces

13. Intermolecular Forces • Dipole-dipole forces are much weaker than covalent or ionic bonds. • Dipole-dipole forces become weaker as the distance between the dipoles increases. In the gas phase, where molecules are far apart, these forces are unimportant. • Strong dipole-dipole forces occur between molecules where hydrogen is bound to a highly electronegative atom, such as nitrogen, oxygen or fluorine. • These bonds are strong because of the great polarity of the bond and the close approach of the dipoles, made possible because of the small size of the hydrogen atom. • These dipole-dipole attractions are called hydrogen bonding.

14. Intermolecular Forces

15. Intermolecular Forces • Hydrogen bonding has an important effect on various physical properties. • The boiling points for the covalent compounds of hydrogen with elements of Group 16 indicate a boiling point of water much higher than expected.

16. Intermolecular Forces

17. Intermolecular Forces • The large electronegativity value of oxygen (2.5) causes the O-H bond (1.4) to be more polar than the S-H (0.4), Se-H (0.3) and Te-H (0.0) bonds. This leads to very strong hydrogen bonding forces among the water molecules. • An unusually large quantity of energy is required to overcome these forces and separate the molecules to produce the gaseous state. Water molecules tend to stay together, even at higher temperatures

18. Intermolecular Forces • Even molecules without dipole moments must exert forces on each other. They exist among noble gas atoms and non-polar molecules. • These forces are called London dispersion forces. • Although we assume the electrons are evenly distributed about the nucleus, it is not true at every instant. • Atoms can develop a temporary dipolar arrangement as the electrons move around the nucleus. This instantaneous dipole can then induce a similar dipole in a neighboring atom. The interatomic attraction formed is weak and short-lived. • London forces become more significant as the size of the atoms or molecules increases.

19. Intermolecular Forces

20. Evaporation and Vapor Pressure • Evaporation, or vaporization, is a process by which the molecules of a liquid escape the liquid’s surface and form a gas. • Evaporation requires energy to overcome the relatively strong intermolecular forces in the liquid. • Because of the strong hydrogen bonding among its molecules in the liquid state, water has a large heat of vaporization. • Practically, this allows water to be used as a coolant.

21. Evaporation and Vapor Pressure • When a given amount of liquid is placed in a closed container the amount of liquid decreases slightly at first and then becomes constant. • The decrease occurs because there is a transfer of molecules from the liquid to the vapor phase. • As the number of vapor molecules increases, it is more likely that some will return to the liquid phase. • The process by which vapor molecules forma liquid is called condensation. • Eventually the same number of molecules leave the liquid state as return to it and the system is in equilibrium.

22. Evaporation and Vapor Pressure

23. Evaporation and Vapor Pressure • The pressure of the vapor present at equilibrium is called the equilibrium vapor pressure or, more commonly, the vapor pressure of the liquid • The vapor pressures of liquids vary widely. • Liquids with high vapor pressures are said to be volatile, they evaporate quickly. • The vapor pressure of a liquid at a given temperature is determined by the intermolecular forces that act among the molecules. • Liquids in which the intermolecular forces are large have relatively low vapor pressures because such molecules need high energies to escape to the vapor phase.

24. Evaporation and Vapor Pressure Which will have a higher vapor pressure, water or methanol (CH3OH)? Water contains two polar O–H bonds while methanol only has one. Therefore, the hydrogen bonding in water is greater than in methanol. Water will have a lower vapor pressure.

25. Evaporation and Vapor Pressure Which will have a higher vapor pressure, methanol (CH3OH) or butanol (CH3-CH2-CH2-CH2-OH)? Each of these molecules has one O–H bond. Butanol is a larger molecule than methanol and has greater London forces, so it is less likely to escape from its liquid. Butanlo has a lower vapor pressure than methanol.

26. The Solid State • Many substances form crystalline solids, with a regular arrangement of their components.

27. The Solid State • There are many different types of crystalline solids. • Sugar and salt both form crystals. Although they both dissolve in water, the properties of the solution are very different. • Salt forms an ionic solid that will conduct electricity in solution due to the Na+ and Cl- ions being free to move around. • Sugar forms a molecular solid that will not conduct electricity in solution. It is composed of neutral molecules, not ions. • Atomic solids are composed of atoms one a single element covalently bonded to each other. Graphite and diamond are forms of carbon in this category.

28. The Solid State

29. The Solid State

30. Bonding in Solids Ionic Solids • Stable substances with high melting point • Held together by strong forces that exist between oppositely charge ions • Structure can be viewed as efficiently packed spheres

31. Bonding in Solids Molecular Solids • Fundamental particle is a molecule • Melt at relatively low temperatures (ice, dry ice) because intermolecular forces are relatively weak • If the molecule has a dipole moment, dipole-dipole forces hold the solid together • If the molecules are non-polar, London dispersion forces hold the solid together • In phosphorous, the distances between the P atoms is smaller than the distance between the P4 molecules because the covalent bond between atoms in the molecule are stronger than the London dispersion forces between molecules. See Fig 13.17

32. Bonding in Solids

33. Bonding in Solids Atomic Solids • Properties vary widely • Solids of noble gases cannot form covalent bonds so they are held together by weak London dispersion forces • Carbon in diamond form has extremely high melting point due to strong covalent carbon-carbon bonds, which lead to a giant molecule

34. Bonding in Metals • Because of the nature of the metallic crystal, other elements can be introduced relatively easily to produce substances called alloys, which contain a mixture of elements and has metallic properties • In a substitutional alloy, some of the host metal atoms are replace by other metal atoms of a similar size

35. Bonding in Metals • An interstitial alloy is formed when some of the interstices (holes) among the closely packed metal atoms are occupied by atoms much smaller than the host atom