Liquids and Solids

1. Liquids and Solids Ch 10.2,10.3 & 10.4 Pg. 353 # 4, 5, 7, 8, 10-14,17, 20-22, 27, 28, 33

2. Liquids exist in the smallest temperature range, so liquids are the least common state of matter . . . • Kinetic Theory Description of the Liquid State • According to the kinetic theory, motion of liquid particles can be described as . . . • a form of matter that has a definite volume and takes the shape of its container. • Kinetic-Theory Description of the Liquid State • Particles in a liquid are in  constant motion however, the particles in a liquid are closer together than the particles in a gas are. • Therefore, the attractive forces between particles in a liquid are more effective than those between particles in a gas. • This attraction is caused by intermolecular forces. • According to the kinetic-molecular theory of liquids, the particles are not bound together in fixed positions.

3. Definite Volume - fixed. Does not vary. They cannot expand to fill a container Fluidity – ability to flow and take shape of container Relative High Density- close arrangement of particles (compared to a gas) making mass/volume ratio higher. Incompressible – much less compressible than gases b/c particles are closer together Dissolving Ability- liquids can dissolve solids, liquids, and gases Ability to Diffuse – mix with other liquids due to constant motion of particles Surface Tension –results from attractive forces between particles on a liquid’s surface. a force that pulls adjacent parts of the liquid’s surface together, thereby decreasing surface area Tendency to Evaporate and Boil - Vaporization is the liquid to gas phase change Tendency to Solidify - Freezing is the liquid to solid phase change Properties of Liquids and the Particles Model – define each property Properties of Fluids

4. 10.2 Questions • Why are liquids more dense than gases? • Molecules are closer together so more molecules in a given area • Why are liquids harder to compress than gases? • Same as above – molecules are closer • Why do liquids diffuse slower than gases? • Particles are not moving as fast as gases • Can a liquid boil without increasing the temperature? How? • Yes – lower the atmospheric pressure

5. 10.3 Solids • “Solid as a rock, “ is the description of solid – something that is hard, unyielding, with a definite shape and volume. Many things other than rocks are solids. In fact, solids are more common than liquids. This diagram shows the particles of a gas, liquid and solid.

6. Kinetic-Theory Description of the Solid StateAccording to the kinetic theory, the motion of solid particles can be described as…. • Lower kinetic energy, less motion, more packed particles, and higher intermolecular forces (IMF) • Intermolecular forces between particles are therefore much more effective in solids. • These hold particles of a solid in relatively fixed positions, with only vibrational movement. • Solids are more ordered than liquids and gases.

7. Properties of Solids and the Particle Model – • Definite shape and volume - solids maintain a definite shape without a container. Volume is constant due to closely packed particles. • Non-fluid- particles can’t flow because particles are held in relatively fixed positions. • Definite melting point - The temperature at which the kinetic energy of the particles are able to overcome the attractive forces holding them together in fixed positions (crystalline only) • High Density- solids are packed more closely than that of a liquid or gas. • Incompressible - particles are packed so close together there is virtually no space between them • Slow Diffusion – much slower than liquids due to the high IMF’s between particles.

8. Crystalline Solids • Classification of crystals by arrangement and shape • Crystal Lattice (define) - The total 3-D array of points that describe the arrangement of the particles – a collection of unit cells. • The smallest portion of the crystal lattice that reveals the 3-D pattern of the entire lattices is the unit cell.

9. Binding Forces in Crystals Body-centered (ex. Li, K, Cr) Simple

10. Types of Crystals Hexogonal (like oranges in a grocery store); (ex. Zn) Face-centered (ex. Cu, Ag, Au)

11. Binding forces in crystals

12. Amorphous Solids • Rubber, glass, plastics and synthetic fibers are called amorphous solids.

13. “Amorphous,” comes from the Greek for “without a shape.” • Unlike crystals, amorphous solids do not have a regular, natural shape, but instead take on whatever shape imposed on them. • Particle arrangement is not uniform; they are arranged randomly, like particles of a liquid. • Examples of amorphous solids – glass, plastic • Amorphous solids are prepared by rapid cooling of thin film materials.

14. Molecular examples Crystalline vs. Amorphous

15. 10.4 Changes of State

16. Equilibrium • What does equilibrium mean? • It is a dynamic condition in which two opposing changes occur at equal rates in a closed system. • What is a closed system? • A substance in a closed beaker, closed off environment

17. When a liquid changes to a vapor, as in evaporation, it absorbs heat energy and can be shown as: • Open system evaporation – liquid + heat vapor • Closed system evaporation – liquid + heat ↔ vapor • When a vapor condenses, as in condensation, it gives off heat energy and can be shown as: • And condensation – vapor  liquid + heat • The liquid vapor equilibrium can be rewritten as: • liquid + heat↔vapor • “The double yields sign represents a reaction at equilibrium”

18. Le Chatelier’s Principle • What is it? LeChatelier • When a system at equilibrium is disturbed by the application of stress, the system reacts to minimize the stress. • Is temperature an example of stress? • Yes. • What happens when you increase the temperature of a system? Equ. shift from heat • ↓ liquid + increased heat ---> ↑ vapor

19. Le Chatelier’s Principle • What happens when you decrease the temperature of a system? • ↓ vapor ---> ↑ liquid + decreased heat • What factor is controlling the decrease and increase of vapor and liquid? • the temperature (heat)

20. Equilibrium Vapor Pressure of a Liquid • What is it? • At equilibrium, the molecules of a vapor exert a specific pressure on its corresponding liquid.

21. When equilibrium vapor pressure of water is graphed, (draw figure 14 below):

22. The strength of attractive forces is independent of temperature. Higher temperatures with resultant higher kinetic energies make these forces less effective. • Liquid water can exist in equilibrium with water vapor only up to a temperature of 374.1ºC. Later you will learn that neither liquid water nor water vapor can exist at temperatures above 374.1ºC.

24. What is equilibrium called when liquid molecules enter into the gaseous state? • Vaporization • Where does this occur? • On the surface of the liquid = evaporation, throughout liquid = boiling • Equilibrium vapor pressure depends on: • a) temperature and pressure • b) boiling point of a liquid (the type of liquid)

25. If a liquid has high intermolecular forces, then what happens to that liquid’s vapor pressure? Why? • vapor pressure ↓ high IMFs = increase hold on the molecules

26. Boiling. Freezing. Melting • What is boiling? • The conversion of a liquid to a vapor, within the liquid as well as its surface when the equilibrium vapor pressure of the liquid is equal to the atmospheric pressure. • What is the boiling point? • The temperature at which the equilibrium vapor pressure of the liquid is equal to the atmospheric pressure (760 torr). • Boiling happens throughout the liquid…evaporation happens on the surface.

27. What is the molar heat of vaporization? • The amount of heat energy required to vaporize one mole of liquid at its boiling point. • How does a pressure cooker work? • It elevates pressure to raise boiling point and shorten cooking time.

28. Freezing and melting • What is the freezing? • The physical change of a liquid to a solid. • What is melting? • The physical change of a solid to liquid. • What is the molar heat of fusion? • The amount of heat energy required to melt one mole of solid at its melting point.

29. solid + heat  Liquid liquid  solid + heat re-write the equation: solid + heat ↔ liquid heat of fusion

30. Are the freezing points and melting points the same temperature? • Yes • at 0°C H2O with 6kJ is a liquid • at 0°C H2O without 6kJ is a solid

31. Chapter 10 Calculations • Molar heat of Vaporization • The amount of heat energy required to vaporize one mole of liquid at its boiling point. • Joules are the standard unit to measure heat energy. • Molar heat of vaporization for water is 40.79 kJ/mole

32. It also takes energy to melt or boil any substance. The amount of energy required to melt or boil a substance can be expressed by the following equations:  • q = nΔHfusionq = change in energy (J) n = number of moles • q = nΔHvaporizationΔHfusion = the molar heat of fusion (kJ/mol) ΔHvaporization = the molar heat of vaporization (kJ/mol) • ΔHfusion and ΔHfusion are constants and correspond to the amount of energy it takes to freeze (fuse) or boil (vaporize) one mole of a substance.

33. Ex1: How much heat energy would be required to vaporize 5.00 moles of H2O • q = ΔHvap·(mol) = = 204 kJ or 204,000 J • Ex2: to vaporize 45.0g of H2O • q = ΔHvap·(mol) = = 102 kJ or 102,000 J

34. when....a liquid evaporates, it absorbs energy. Energy is used to overcome attractive forces.The energy doesn’t increase the average energy of the particles, so the temperature doesn’t change. • when...a liquid evaporates, it takes energy from its surroundings that’s why alcohol feels cool to the skin. • it’s also why we get cold when getting out of the shower

35. Molar Heat of Fusion • The amount of heat energy required to melt one mole of a solid at its melting point. • The molar heat of fusion of water is 6.008 kJ/mole. • Ex1: How much energy would be required to melt 12.75 moles of ice? • q = ΔHfus·(mol) = = 76.60 kJ

36. Ex2: to melt 6.48 x 1020 kg of ice?

37. Ex3: - How much ice can be melted by 2.9 x 104 J?

38. Heat and Temperature – there is a difference • Heat is the amount of energy a chemical has, frequently measured in joules (J). Because we can’t directly measure heat, we have to measure “temperature”, which reflects how much kinetic energy an object has (as measured in °C or Kelvins).

39. Heat and Temperature – there is a difference • Heat transfers between objects – flows from hot to cold - Law of Conservation of Energy • Ex1:ice cube in a thermos of hot water - ice melts, water cools - same amount of heat • SI unit of heat - Joule (J) -calorie is also used frequently • Calorie - the amount of energy required to raise the temperature of 1 g of water by 1 oC • (Calories – capital letter – really means kilocalories – used in food energy measurement) 1.000 calorie = 4.184 Joules This is not in your notes. Just make sure you write down the above conversion.

40. Three factors affect how much heat an object absorbs or loses • Mass of the object • Change in temperature • final temperature - initial temperature • if there is no change in temperature, no heat flows • Specific Heat • specific heat (Cp): heat required to raise the temp. of 1 g of material by 1 K or 0c • different materials have different specific heats

41. Specific Heat • Ex, which would you rather use to pull a pan from a hot oven- an oven mitt or a sheet of aluminum foil? The aluminum foil will transmit the heat easily while the oven mitt is a much better insulator. The reason: Oven mitts have a higher heat capacity (specific heat) than aluminum.

42. Computing heat to determine how much heat is required to heat a material • It takes energy to make the temperature of anything increase. The relationship between energy and temperature is shown by the equation: • q = m cpΔT • q = heat/energy (Joules or calories) • m = mass - grams or moles • Cp = specific heat • ΔT = Tfinal - Tinitial

43. Specific Heat Problems • For water, Cp = 1.000 cal/goC or 4.184 J/goC • Ex1:How much heat is required to heat 75 g of Iron (Cp = 0.444 J/gCo) from 15.5 to 57.0oC? • EX2: How many calories does it take to heat 20. g of water from 10.0 to 40.0oC? Also how many J?

44. Specific Heat Problems- • Ex3: What is the specific heat of an object if 250 calories will heat 55 g of it from 25 to 100.0 oC? • Ex4: - If a 100.0 g sample of silver (Cp = .237 J/g oC) at 80.0 Co loses 50. calories, what will its final temperature be?

45. Temperature and Phase Changes Flat sections at boiling/melting why? All energy input is directed at changing phase, so there is no increase in temp.

46. Temperature and Phase Changes • 3 formulas to use: • q = mCpt for sections A, C, E • q = mHfus for section B • q = mHvap for section D

47. Temperature and Phase Change • It is usually assumed that more heat means higher temperature, but not when changing phase. • EX:1 If a sample of water at 20.0 oC is heated by a hot plate that gives off 250.0 J, how grams of water are in the sample if the temperature rises to 30.0 oC?

48. Temperature and Phase Change • EX 2: How many kJ are needed to convert 25.0 g of water from a liquid at 50.0° C to a gas to 100.0° C.

49. Temperature and Phase Change • EX 3: If 75 g H2O is at -5.0oC and is heated to 115.0oC. How much total heat in calories is required? (Convert to joules after)