How to Use This Presentation

1. How to Use This Presentation • To View the presentation as a slideshow with effects select “View” on the menu bar and click on “Slide Show.” • To advance through the presentation, click the right-arrow key or the space bar. • From the resources slide, click on any resource to see a presentation for that resource. • From the Chapter menu screen click on any lesson to go directly to that lesson’s presentation. • You may exit the slide show at any time by pressing the Esc key.

2. Resources Chapter Presentation Bellringer Transparencies Sample Problems Visual Concepts Standardized Test Prep

3. States of Matter and Intermolecular Forces Chapter 11 Table of Contents Section 1States and State Changes Section 2Intermolecular Forces Section 3Energy of State Changes Section 4Phase Equilibrium

4. Section1 States and State Changes Chapter 11 Bellringer • Define the term surface tension in your own words. • What state of matter is associated with surface tension?

5. Section1 States and State Changes Chapter 11 Objectives • Relate the properties of a state to the energy content and particle arrangement of that state of matter. • Explain forces and energy changes involved in changes of state.

6. Section1 States and State Changes Chapter 11 States of Matter • Most substances can be in three states: solid, liquid, and gas. • Solid Particles Have Fixed Positions • The particles in a solid are very close together and have an orderly, fixed arrangement. • Solid particles can vibrate only in place and do not break away from their fixed positions. • Solids have fixed volumes and shapes.

7. Mercury in Three States Section1 States and State Changes Chapter 11

8. Section1 States and State Changes Chapter 11 States of Matter, continued Liquid Particles Can Move Easily Past One Another • The particles in a liquid are very close together andhave a random arrangement. • Liquid particles have enough energy to be able to move past each other readily, which allows liquids to flow. • Liquids have fixed volumes but can flow to take the shape of the lower part of a container.

9. Chapter 11 Section1 States and State Changes States of Matter, continued Liquid Forces Lead to Surface Wetting and Capillary Action • Liquid particles can have cohesion, attraction for each other. • Liquid particles can also have adhesion, attraction for particles of solid surfaces. • The balance of cohesion and adhesion determines whether a liquid will wet a solid surface. • The forces of adhesion and cohesion will pull water up a narrow glass tube, called a capillary tube.

10. Chapter 11 Section1 States and State Changes Comparing Cohesion and Adhesion

11. Chapter 11 Section1 States and State Changes States of Matter, continued Liquids Have Surface Tension • Below the surface of a liquid, the particles are pulled equally in all directions by cohesive forces. • However, surface particles are pulled only sideways and downward, so they have a net downward force. • It takes energy to oppose this net force and increase the surface area. • The tendency of liquids to decrease surface area to the smallest size possible is surface tension.

12. Chapter 11 Section1 States and State Changes Surface Tension

13. Chapter 11 Section1 States and State Changes States of Matter, continued Gas Particles Are Essentially Independent • The particles in a gas are very far apart and have a random arrangement. • The attractive forces between particles in a gas do not have a great effect, so the particles move almost independently of one another. • The shape, volume, and density of an amount of gas change depending on the size and shape of the container.

14. Chapter 11 Section1 States and State Changes Solid, Liquid and Gas

15. Chapter 11 Section1 States and State Changes Changing States • Most substances can undergo six changes of state: freezing, melting, evaporation, condensation, sublimation, and deposition. • Temperature, Energy, and State • Generally, adding energy to a substance will increase the substance’s temperature. • But after a certain point, adding more energy will cause a substance to experience a change of state instead of a temperature increase.

16. Changes of State Chapter 11 Section1 States and State Changes

17. Chapter 11 Section1 States and State Changes Changing States, continued Liquid Evaporates to Gas • Energy is required to separate liquid particles. They gain energy when they collide with each other. • If a particle gains a large amount of energy, it can leave the liquid’s surface and join gas particles. • Evaporation is the change of state from liquid to gas. Evaporation is an endothermic process. • Boiling point is the temperature and pressure at which a liquid and a gas are in equilibrium.

18. Chapter 11 Section1 States and State Changes Changing States, continued Gas Condenses to Liquid • When gas particles no longer have enough energy to overcome the attractive forces between them, they go into the liquid state. • Condensation is the change of state from a gas to a liquid. Condensation is an exothermic process. • Condensation can take place on a cool night, causing water vapor in the air to form dew on plants.

19. Visual Concepts Chapter 11 Vaporization and Condensation

20. Chapter 11 Section1 States and State Changes Changing States, continued Solid Melts to Liquid • As a solid is heated, the particles vibrate faster and faster in their fixed positions. • At a certain temperature, some of the molecules have enough energy to break out of their fixed positions. • Melting is the change of state from solid to liquid. Melting is an endothermic process. • Melting point is the temperature and pressure at which a solid becomes a liquid.

21. Chapter 11 Section1 States and State Changes Changing States, continued Liquid Freezes to Solid • As a liquid is cooled, the movement of particles becomes slower and slower. • At a certain temperature,the particles are pulled together into the fixed positions of the solid state. • Freezing is the change of state from a liquid to a solid. Freezing is an exothermic process. • Freezing point is the temperature at which a substance freezes.

22. Visual Concepts Chapter 11 Freezing

23. Chapter 11 Section1 States and State Changes Changing States, continued Solid Sublimes to Gas • The particles in a solid are constantly vibrating. Some particles have higher energy than others. • Particles with high enough energy can escape from the solid. • Sublimation is the change of state from solid to gas. • Sublimation is an endothermic process.

24. Chapter 11 Section1 States and State Changes Changing States, continued Gas Deposits to Solid • Molecules in the gaseous state become part of the surface of a crystal. • When a substance changes state from a gas to a solid, the change is often called deposition. • Deposition is an exothermic process.

25. Visual Concepts Chapter 11 Comparing Sublimation and Deposition

26. 1. Describe what happens to the shape and volume of a solid, a liquid, and a gas when you place each into separate, closed containers. The solid will keep its shape & volume. The liquid will keep the same volume but will take the shape of its container. The gas will change in volume & shape by filling the entire container. 2. What is surface tension? Is the force that acts on the surface of a liquid & that tends to minimize the area of the surface. 3. You heat a piece of iron from 200 to 400 K. What happens to the atoms’ energy of random motion? The energy of random motion doubles. Section 11.1 Review, pg. 384

27. 4. When water boils, bubbles form at the base of the container. What gas has formed? Water vapor 5. What two terms are used to describe the temperature at which solids and liquids of the same substance exist at the same time? Melting & freezing point 6. How are sublimation and evaporation similar? In both processes, molecules leave a surface and go into the gaseous phase. 7. Describe an example of deposition. Frost forming on a surface on a cold night.

28. 8. The densities of the liquid and solid states of a substance are often similar. Explain. In both liquids and solids, the particles are nearly in contact with one another. 9. How could you demonstrate evaporation? Hanging a wet cloth up to dry. 10. How could you demonstrate boiling point? Putting water in a kettle and heating the water until bubbles form steadily and the temperature remains constant.

29. 11. You are boiling potatoes on a gas stove, and your friend suggests turning up the heat to cook them faster. Will this idea work? No, the water is already at the boiling pt., so increasing the heat will not cause the temperature to Rise. It will only increase the rate of evaporation. 12. A dehumidifier takes water vapor from the air by passing the moist air over a set of cold coils to perform a state change. How does a dehumidifier work? Water vapor is removed from the air when it condenses onto the cold coil.

30. 13. Water at 50°C is cooled to −10°C. Describe what will happen. When the water reaches 0 C, ice will begins to form. When all of the water is solid , the ice will cool to -10 C 14. How could you demonstrate melting point? Put some ice cube and water in a glass. One some of the cubes have melted, you would determine the temperature. 15. Explain why changes of state are considered physical transitions and not chemical processes. They don’t involve a change in the identity of the substance.

31. Section2 Intermolecular Forces Chapter 11 Bellringer • List any terms you know that use the prefixes inter- and intra-.

32. Section2 Intermolecular Forces Chapter 11 Objectives • Contrast ionic and molecular substances in terms of their physical characteristics and the types of forces that govern their behavior. • Describe dipole-dipole forces. • Explain how a hydrogen bond is different from other dipole-dipole forces and how it is responsible for many of water’s properties.

33. Section2 Intermolecular Forces Chapter 11 Objectives, continued • Describe London dispersion forces, and relate their strength to other types of attractions.

34. Section2 Intermolecular Forces Chapter 11 Comparing Ionic and Covalent Compounds • It takes energy to overcome the forces holding particles together. • Thus, it takes energy to cause a substance to go from the liquid to the gaseous state. • The boiling point of a substance is therefore a good measure of the strength of the forces that hold the particles together. • Melting point also relates to attractive forces between particles.

35. Section2 Intermolecular Forces Chapter 11 Comparing Ionic and Covalent Compounds, continued • Most covalent compounds melt at lower temperatures than ionic compounds do.

36. Section2 Intermolecular Forces Chapter 11 Comparing Ionic and Covalent Compounds, continued Oppositely Charged Ions Attract Each Other • Ionic substances generally have much higher forces of attraction than covalent substances. • For small ions, attractions between ions of opposite charge hold the ions tightly in a crystal lattice. • These attractions are overcome only by heating to very high temperatures.

37. Section2 Intermolecular Forces Chapter 11 Comparing Ionic and Covalent Compounds, continued Oppositely Charged Ions Attract Each Other, continued • If the ions are larger, then the distances between them are larger and the forces are weaker. • Thus, ionic compounds with small ions have high melting points.

38. Section2 Intermolecular Forces Chapter 11 Comparing Ionic and Covalent Compounds, continued Intermolecular Forces Attract Molecules to Each Other • Intermolecular forces are the forces of attraction between molecules of covalent compounds. • Intermolecular forces include dipole-dipole forces and London dispersion forces.

39. Section2 Intermolecular Forces Chapter 11 Dipole-Dipole Forces Dipole-Dipole Forces Affect Melting and Boiling Points • Dipole-dipole forces are interactions between polar molecules. • When molecules are very polar, the dipole-dipole forces are very significant. • The more polar the molecules are, the higher the boiling point of the substance.

40. Section2 Intermolecular Forces Chapter 11 Dipole-Dipole Forces

41. Section2 Intermolecular Forces Chapter 11 Hydrogen Bonds • A hydrogen bond is a dipole-dipole force occurring when a hydrogen atom that is bonded to a highly electronegative atom of one molecule is attracted to two unshared electrons of another molecule. • In general, compounds with hydrogen bonding have higher boiling points than comparable compounds.

42. Section2 Intermolecular Forces Chapter 11 Hydrogen Bonds, continued • As the electronegativity difference of the hydrogen halides increases, the boiling point increases. • The boiling points increase somewhat from HCl to HBr to HI but increase a lot more for HF due to the hydrogen bonding between HF molecules.

43. Section2 Intermolecular Forces Chapter 11 Hydrogen Bonds, continued Hydrogen Bonds Form with Electronegative Atoms • Strong hydrogen bonds can form with a hydrogen atom that is covalently bonded to very electronegative atoms in the upper-right part of the periodic table: nitrogen, oxygen, and fluorine. Hydrogen Bonds Are Strong Dipole-Dipole Forces • The combination of the large electronegativity difference (high polarity) and hydrogen’s small size accounts for the strength of the hydrogen bond.

44. Section2 Intermolecular Forces Chapter 11 Hydrogen Bonding

45. Section2 Intermolecular Forces Chapter 11 Hydrogen Bonds, continued Hydrogen Bonding Explains Water’s Unique Properties • Each water molecule forms multiple hydrogen bonds, so the intermolecular forces in water are strong. • The angle between the two H atoms is 104.5°. • When water forms ice, the ice crystals have large amounts of open space. Thus, ice has a low density. • Water is unusual in that its liquid form is denser than its solid form.

46. Ice and Water Section2 Intermolecular Forces Chapter 11

47. Section2 Intermolecular Forces Chapter 11 London Dispersion Forces • A substance with weak attractive forces will be a gas because there is not enough attractive force to hold molecules together as a liquid or a solid. • However, many nonpolar substances are liquids. • What forces of attraction hold together nonpolar molecules and atoms? • London dispersion forces are the dipole-dipole force resulting from the uneven distribution of electrons and the creation of temporary dipoles.

48. Section2 Intermolecular Forces Chapter 11 London Dispersion Force

49. Section2 Intermolecular Forces Chapter 11 London Dispersion Forces, continued London Dispersion Forces Exist Between Nonpolar Molecules • As molar mass increases, so does the number of electrons in a molecule. • London dispersion forces are roughly proportional to the number of electrons present. • Thus, the strength of London dispersion forces between nonpolar particles increases as the molar mass of the particles increases.

50. Section2 Intermolecular Forces Chapter 11 London Dispersion Forces, continued London Dispersion Forces Result from Temporary Dipoles • The electrons in atoms can move about in orbitals and from one side of an atom to the other. • When the electrons move toward one side of an atom or molecule, that side becomes momentarily negative and the other side becomes momentarily positive. • When the positive side of a momentarily charged molecule moves near another molecule, it can attract the electrons in the other molecule.