Chapter 13 States of Matter

1. Chapter 13 States of Matter

2. FOOD FOR THOUGHT… • What is the relationship between solids, liquids, and gases? How are they the same? How are they different? • Make a table or list answering the above question. You will need your answers tomorrow.

3. Flemish physician Jan Baptista Van Helmont used the Greek word chaos, meaning without order, to describe the products or reactions that had no fixed shape or volume. From chaos came the term gas. 13.1 - Gases

4. Around 1860, Ludwig Boltzmann and James Maxwell each proposed a model to explain the properties of gases. The model is known as the kinetic-molecular theory. Kinetic-molecular theory describes the behavior of gases in terms of particles in motion. 13.1 - Gases

5. Kinetic – molecular theory: (pages 386, 419-420) Gas particles do not attract or repel each other. Why? Gas particles are much smaller than the distance between them. The theory assumes the particles have virtually no volume. What process can be applied to gases because of the very small particle size? Gas particles are in constant, random motion. Because of this, gas particles spread out and mix. How do gas particles move? What may cause their motion to change? 13.1 - Gases

6. Kinetic – molecular theory: (pages 386, 419-420) No kinetic energy is lost when gas particles collide with each other or with the walls of their container. Why? How can you describe elastic collisions? What can be said about the total kinetic energy of a container of gas? All gases have the same average kinetic energy at a given temperature. What happens to total energy of a gas as temperature increases? What happens to total energy of a gas as temperature decreases? What two factors determine kinetic energy? What is temperature? What do all gases have in common? 13.1 - Gases

7. Food for thought… Will the kinetic-molecular theory work for liquids and solids? Explain your reasoning.

8. How does kinetic-molecular theory explain the behavior of gases? Density The idea that there is a lot space between particles explains the low density (mass / volume) that gases have. There are fewer molecules of gas in a given amount of volume. Compression and Expansion Gas molecules can be compressed (pressed into a smaller space using a piston). Gas molecules can expand as pressure from the piston is released. What happens to the density as a gas is compressed? What happens to the density as a gas expands? 13.1 - Gases

9. How does kinetic-molecular theory explain the behavior of gases? Diffusion The idea that there are no significant forces of attraction between gas particles supports the fact that gases can flow easily past each other. When gases mix, what will eventually happen to their concentration? What is diffusion? In which direction does diffusion occur? What does the rate of diffusion depend upon? Effusion What is effusion? How can effusion be compared to diffusion? 13.1 - Gases

10. How does kinetic-molecular theory explain the behavior of gases? Effusion Graham’s law of effusion – the rate of effusion for a gas is inversely proportional to the square root of its molar mass Graham’s law applies to diffusion rates also. Rate of effusion 1 1 SQRT(molar mass) Rate A = SQRT(molar mass B / molar mass A) Rate B 13.1 - Gases

11. Food for thought… • What state of matter is He? • What is happening to the He atoms inside a balloon to keep the balloon inflated? Use the kinetic-molecular theory to support your answers.

12. 13.1 - Gases TIME TO WORK! Effusion / diffusion problems (1-2), page 388 Dalton’s Law problems (4-6), page 392 Problem-solving lab, page 390 Calculate the ratio of effusion rates for pairs of the noble gases.

13. Intramolecular Forces The forces that hold particles together in ionic, covalent, and metallic bonds are called intramolecular forces. These are forces occurring within the chemical compound. Intermolecular Forces Intermolecular forces happen between or among like molecules of a substance. There are 3 intermolecular forces we will discuss: dispersion forces, dipole-dipole forces, and hydrogen bonds. 13.2 – Forces of Attraction

14. Intramolecular Forces All intermolecular forces are weaker than intramolecular bonding forces. 13.2 – Forces of Attraction • Relative Strength of Molecular Forces • Covalent network > ionic bonds > metallic bonds > • Hydrogen bonds > dipole-dipole forces > London dispersion • forces

15. Food for thought… • Many words can be understood by looking at their parts. • Match the following word parts to their meanings. -ion therm- -ic -ize endo- exo- Heat The result of an action Related to To become Inside Outside

16. London dispersion forces – are weak forces that result from temporary shifts in the density of electrons in electron clouds (draw sketch page 394) are named after the German-American physicist who first described them, Fritz London are weak forces because they are based on temporary dipoles are the dominant force of attraction between identical nonpolar molecules dispersion force strength: I > Br > Cl > F explains why F and Cl are gases, Br is liquid, and I is solid at room temperature 13.2 – Forces of Attraction

17. Dipole-dipole forces – are forces of attraction between oppositely charged regions of polar molecules. are weak forces of attraction that result from permanent dipoles within polar molecules Some regions of a polar molecule are always partially negative and other regions are always partially positive. Neighboring polar molecules orient themselves so that oppositely charged regions line up. (sketch diagram page 394) The degree of polarity in a molecule depends on the relative electronegativity values of the elements in the molecule. Dipole forces are stronger than dispersion forces as long as the molecules being compared have about the same mass. 13.2 – Forces of Attraction

18. Bell ringer… • Which are stronger, intermolecular or intramolecular forces? • Which specific force is strongest? • Which is weakest?

19. Hydrogen bonds – are a special type of dipole attractive force between highly polar molecules. are dipole-dipole attractions that occur between molecules containing a hydrogen atom bonded to a small, highly electronegative atom with at least one lone electron pair. (sketch diagram page 395) Hydrogen must be bonded to either a F, O, or N atom because these atoms are electronegative enough to cause a large partial positive charge on the hydrogen, yet are small enough that their lone pairs of electrons can come close to hydrogen atoms. In hydrogen bonds, hydrogen atoms on one molecule of a substance are attracted to the negative end of another molecule of the substance. Hydrogen bonds explain why water is liquid at room temp. 13.2 – Forces of Attraction

20. Although kinetic-molecular theory was developed to explain the behavior of gases, the model can be applied to liquids and solids. However, you must consider the forces of attraction between particles as well as their energy of motion. Why consider the force of attraction between particles of liquids and solids? 13.3 – Solids and Liquids In liquids and solids, particles are closer together because of stronger forces of attraction between them. In liquids, the forces of attraction limit their range of motion so that particles are closely packed in a fixed volume. In solids, the forces are so strong that motion is limited to vibrations around a fixed location. Therefore, solids have more order in their particles.

21. Food for thought… • Put the following attractive forces in the proper category: intermolecular or intramolecular dispersion forces dipole-dipole covalent bonds metallic bonds hydrogen bonds ionic bonds • Put the attractive forces in order from weakest to strongest.

22. Liquids - define the following terms related to properties of liquids Density Compression Fluidity Viscosity Temperature (relationship to viscosity) Surface tension (define and give an example) Surfactants Capillary action (define and give an example) 13.3 – Solids and Liquids

23. Solids - define the following terms related to properties of solids Density Compressibility Fluidity Crystalline solid Unit cell Molecular solids Covalent network solids Ionic solids Metallic solids Amorphous solids 13.3 – Solids and Liquids

24. 13.3 – Solids and Liquids • Sketch and label 3 basic solid crystal lattice structures (page 400).

25. 6 possible transitions between phases 13.4 – Phase Changes GAS SUBLIMATION CONDENSATION VAPORIZATION DEPOSITION LIQUID SOLID MELTING FREEZING

26. Phase changes requiring energy input Melting What is melting? What is heat? Describe what happens when ice melts. Why is the energy required to melt salt much greater than the energy to melt ice? What is melting point? 13.4 – ENDOTHERMIC CHANGES

27. Phase changes requiring energy input Melting Endothermic Phase Changes Melting is when energy added to a solid is great enough to break the forces holding the molecules together. The solid phase changes to a liquid. Melting point of a solid is the temperature at which forces holding its structure together are broken and it becomes a liquid. Heat is the transfer of energy from a higher temperature to a lower temperature. When ice melts, molecules on the surface absorb enough Energy to break the hydrogen bonds. The molecules move apart and become a liquid. The ionic bonds in salt (sodium chloride) are much stronger than the hydrogen bonds in ice.

28. Phase changes requiring energy input Vaporization When can the temperature of a melting substance start to rise? How much energy does a substance need to vaporize? What is a vapor? What is vaporization? What is evaporation? How long does it take for evaporation to occur? 13.4 – ENDOTHERMIC CHANGES

29. Phase changes requiring energy input Vaporization Endothermic Phase Changes The temperature of a melting substance will start to rise after all the solid substance has melted. Evaporation is vaporization that occurs only at the surface of a liquid. To vaporize, a substance needs enough energy to overcome the forces of attraction holding the molecules together in the liquid. The time for evaporation depends upon the amount of liquid and the amount of energy available. A vapor is a substance that is ordinarily a liquid at room temperature. Vaporization is the process by which a liquid changes to a gas or vapor.

30. Phase changes requiring energy input Vaporization What is vapor pressure? How does a rise in temperature affect vapor pressure? What is the boiling point? As temperature increases, what happens to kinetic energy of molecules? 13.4 – ENDOTHERMIC CHANGES

31. Phase changes requiring energy input Vaporization Endothermic Phase Changes Vapor pressure is the pressure exerted by a vapor over a liquid. Vapor pressure increases with increasing temperature. The boiling point is the temperature at which the vapor pressure of a liquid equals the outside or atmospheric pressure. Kinetic energy of molecules increases with temperature.

32. Phase changes requiring energy input Sublimation What is sublimation? Name some substances that sublime. Endothermic Phase Changes Sublimation is the process by which a solid changes directly to a gas without becoming a liquid. Solid iodine, solid carbon dioxide (dry ice), moth balls, air fresheners, ice cubes left in freezer.

33. Phase changes that release energy Condensation Name some examples of condensation. What is condensation? Describe what happens to molecules when they condense. What happens when hydrogen bonds form in liquid water? Name some circumstances that will cause condensation. 13.4 – Exothermic Changes

34. Phase changes that release energy Condensation Exothermic Phase Changes Examples: dew, frost, water on outside of a glass, water on a window pane, clouds, fog, air conditioning condensate Condensation can be caused by vapor molecules contacting a cold surface (drink glass, dew) or cold air (fog, clouds). Condensation is the process by which a gas or a vapor becomes a liquid. It is the reverse of vaporization. When vapor molecules condense, they lose energy, slow down, and form bonds with each other when they collide. The bonded molecules are more dense and become a liquid. When hydrogen bonds are formed in water, energy is released.

35. Phase changes that release energy Deposition What is deposition? Give some examples of deposition. Exothermic Phase Changes Deposition is the process by which a substances changes from a gas or vapor to a solid without first becoming a liquid. It is the reverse of sublimation. Example: snow

36. Phase changes that release energy Freezing What is freezing? Describe what happens when something freezes. What is freezing point? 13.4 – Exothermic Changes

37. Phase changes that release energy Freezing Exothermic Phase Changes Freezing is the phase change of a liquid to a solid. During freezing, heat is removed from a substance, the molecules lose kinetic energy and slow down. When enough energy has been removed, the molecules become fixed in a set position. The freezing point is the temperature at which a liquid is converted into a crystalline solid. The freezing point and melting point temperatures are the same for a substance.

38. 15. SUMMARIZING ________thermic _______  Solid Gas  ______ ______  Solid Order of the molecules is __________________. _______thermic Solid  _______ _______  Gas Liquid  ______ Order of the molecules is __________________.

39. SUMMARIZING ENDOTHERMIC • Molecules are_________ energy during an endothermic phase change. • What kind of energy? • Describe the change in motion of molecules during melting, vaporization, and sublimation. gaining Kinetic energy – the energy of motion The molecules gain energy and move faster and further apart.

40. SUMMARIZING EXOTHERMIC • Molecules are__________ kinetic energy during an exothermic phase change. • Describe the change in motion of molecules during freezing, condensation, and deposition. losing The molecules lose energy and slow down getting closer together.

41. Phase diagrams Temperature and pressure are the two variables that combine to control the phase of a substance. A phase diagram is a graph of pressure versus temperature that shows in which phase a substance exists under different conditions of temperature and pressure. What is the triple point? All 6 phase changes can occur at the triple point. 13.4 – Phase Changes

42. Phase diagrams What is the critical point? 13.4 – Phase Changes

43. Problem solving! Add to your tools! • Read the problem statement thoroughly. • Identify the “givens” and write them down. • Identify the “unknowns” and write them down. • Determine the scientific principles you will use. • Devise a strategy for attacking the problem.

44. Scientific Principles • Name principles you might use to solve various types of physical science problems. • Law of Conservation of Mass • Rules for Balancing Formulas and Equations • Periodic Table • Solving Mathematical Equations, like F = ma, velocity, acceleration, momentum • Rules for Balanced and Unbalanced Forces • Gravity – free fall, weight, acceleration due to gravity • Newton’s Laws

45. Strategies you have used to solve problems… • Using graphs to show data. • Drawing diagrams to show atomic structure. • Using the Periodic Table to determine trends in atomic radius and ionization energy. • Balancing chemical equations. • Using equations to calculate an unknown value. • Drawing a force diagram and showing the magnitude and direction of the forces acting on an object.