Chemistry 104 Chapter Seven

1. Chemistry 104 Chapter Seven Gases, Liquids and Solids A neutron walks into a bar, sits down and asks for a drink. Finishing, the neutron asks "How much?“ The bartender says, "For you, no charge."

2. You will be able to: • Describe kinetic molecular theory of • matter as applied to the 3 states of matter • Differentiate btwn various Gas Laws & be • able to use them to calculate changes • Apply Gas Laws to real life situations • Explain importance of the following • intermolecular attractions: • Dipole-dipole • H bonding: primarily H2O as an example • London Forces

3. The Kinetic Molecular Theory • Kinetic Energy (K.E.) • random motion • Disruptive force - keeps particles independent from each other • A type of energy • can be transferred in collisions • Molecular • matter composed of tiny particles • Definite sizes & characteristics - immutable

4. Potential energy (P.E.) • Energy in matter due to: • its position (gravitational potential E) • its condition • its composition

5. Potential energy (P.E.) • Chemical Potential energy is dependent upon cohesive forces – electrostatic attraction or repulsion btwn particles • P.E. is independent of temperature • K.E. is very dependent upon temperature

6. Physical Properties of Matter D = “change” • Volume & shape • Density • Compressibility: D in volume due to D in pressure • Thermal expansion: D in volume due to temperatureD

7. Solid: particles (atoms, molecules, or ions) close together & vibrate. Liquid: particles still close together, freely slide over one another. Gas: particles in constant random motion, each particle independent of others.

8. Kinetic Energy & States of Matter Solids: PE (cohesive forces) > KE (disruptive forces) Liquids: PE (cohesive forces) ~ = KE (disruptive forces) Gases: KE (disruptive forces) MUCH > PE (cohesive forces)

9. Comparison of States of Matter Solids: • Definite volume / Definite shape • High density = packed together • Thermal expansion is very low • Strong cohesive forces between particles • Vibrations increase • Low compressibility • very little space between particles • Pressure has little effect on volume

10. Liquids: • Definite volume / NO specific shape • High density • still packed together, though usually not as tightly as solids • Thermal expansion is low • Some cohesive forces between particles • Neighboring particles still touching • increase vibrations • Low compressibility • very little space between particles • Pressure has little effect on volume

11. Gases: • Nodefinite volume / No definite shape • Low density = spread out particles • Thermal expansion is moderate • minimal cohesive forces between particles • Particle speed increases / size remains same • Space between particles increases (expands) • Large compressibility • LOTS of space between particles • Pressure has a big effect on volume

12. Compress a gas: empty space in container decreases. Size of molecules does not D. Molecules just move closer together.

13. Gases, Liquids, and SolidsIntroduction Whether a substance is a gas, liquid, or solid depends on balance btwn KEof its particles and strength of interactions btwn the particles.

14. Gases, Liquids, & Solids a The symbol “~” means approximately.

15. A supercritical fluid shows the properties of both a liquid and a gas. Supercritical fluids: Compressed & heated gas. Properties of liquid and gas simultaneously.

16. increasing both T & P: liquid becomes less dense (thermal expansion) & gas becomes more dense as pressure rises. Distinction btwn gas & liquid disappears. Supercritical CO2: used to dissolve materials; replaces organic solvents; generates less waste. (e.g. decaffeinated coffee)

17. History of The Gas Laws • Gas laws started to evolve in 1643 • Barometer revolutionary tool! • gas laws evolved until 1873: van der Waals equation.  • This occurred before modern atomic theory.  • Gas laws led to numerous concepts: • the mole, temperature, formula weight, absolute zero, kinetic energy, & stoichiometric coefficients.

18. Hg is > 13 xas dense asH2O a water barometer would use a tube > 30 ft long First barometer was built in 1643by Galileo's secretary Evangelisto Torricelli

19. Essential components of a mercury barometerare: a graduated glass tube, a glass dish, liquid mercury, & air pressure

20. The Gas Laws Four variables usually define the state (i.e. condition) of a gas: Temperature, T Pressure, P Volume, V # of moles, n (quantity of particles) Gas laws are designed to look for quantitative relationships between T, P, n and V

21. What exactly is “Pressure?” Pressure = Force/Area • Same forceexerted over smaller area creates a greater P than over larger area 2,000 lbs / 144 in2 vs 100 lbs / 1 in2? 13.8 psi 100 psi

22. Pressure = Force/Area • Pressure in a fluid is exerted in all directions • round balloons • Buoyancy force • the unit of pressure "lbs/sq in" is often written as "psi" ("pounds per square inch"). • Other units include: 1 atm = 760 mm Hg = 760 Torr = 14.7 psi

23. kinetic-molecular theory of gases: • Gas particles move randomly & rapidly. • Size of gas particles is small compared to • space btwn particles. • Gas particles exert no attractive forces • on each other. • KE of gas particles increases with • increasing temp. • When gas particles collide with each other, • they rebound & travel in new directions.

24. 7.2B Gas Pressure When gas particles collide with walls of a container, they exert a pressure. Pressure (P) is force (F) exerted per unit area (A). Force F Pressure = = Area A 760. mm Hg 760. Torr 14.7 psi 101,325 Pa ~101 kPa 1 atmosphere (atm) =

25. Chapter 7 Clicker Question #1: Scuba divers typically start with a tank of air compressed to 3,000 psi. What is the pressure in atmospheres? 204 atm 102 atm 44,100 atm 51.7 atm 3,000 psi x 1 atm = 204 atm 14.7 psi

26. Figure 7.3 Systolic = max P in artery when heart contracts Diastolic = minimum P when heart relaxes 100 to 120 systolic over 60m to 80 for diastolic

27. Chapter 7 Clicker Question #2: A person with 120 over 80 blood pressure is measuring pressure in mm of Hg. What is this pressure in atmospheres? A. 0.158 & 0.105 atm B. 91,200 & 60,800 atm C. 6.16 & 5.44 atm D. 1,764 & 1,176 atm 120 mm of Hg x 1 atm = 0.158 atm 760 mm of Hg 80 mm of Hg x 1 atm = 0.105 atm 760 mm of Hg

28. 7.3A Boyle’s Law: How Pressure & Volume of a Gas Are Related • For a fixed amount of gas • at constant temp, • pressure and volume of the gas • are inversely related. • If one increases, the other decreases. • Product of the two quantities is a constant, k • Pressure x Volume = constant • P x V = k

29. 1662 Boyle's law: At constant temperature P and V are inverseproportions

30. P1V1 = P2V2 • Suppose Volume is increased: • Gas molecules - farther to go & hit container walls less often. • Gas Pressure is less: fewer molecule impacts per unit time. • If Volume is decreased: • Gas molecules - shorter distance to go, striking walls more often. • Pressure increases; more molecule impacts per unit time.

31. Boyle’s Law If volume of cylinder of gas is halved, pressure of gas inside cylinder doubles. This behavior can be explained by the equation: P1V1 = P2V2 initial conditionsnew conditions

32. Boyle’s Law HOW TO Use Boyle’s Law to Calculate a New Gas Volume or Pressure If a 4.0-L container of helium gas has a pressure of 10.0 atm, what pressure does the gas exert if the volume is increased to 6.0 L? Example Identify the known quantities and the desired quantity. Step 1 P1 = 10.0 atm V1 = 4.0 L V2 = 6.0 L P2 = ? known quantities desired quantity

33. Boyle’s Law HOW TO Use Boyle’s Law to Calculate a New Gas Volume or Pressure Write the equation and rearrange it to isolate the desired quantity on one side. Step 2 Solve for P2 by dividing both sides by V2. P1V1 = P2V2 P1V1 = P2 V2

34. Boyle’s Law HOW TO Use Boyle’s Law to Calculate a New Gas Volume or Pressure Step 3 Solve the problem. P1V1 (10.0 atm)(4.0 L) P2 = = = 6.7 atm (6.0 L) V2 Answer Liters cancel.

35. Boyle’s Law and Breathing • To inhale: • Rib cage expands • and diaphragm lowers. • This increases • volumeof lungs. • Increasing volume • causes pressureto • decrease. • Air is drawn into lungs • to equalize pressure.

36. To exhale: • Rib cage contracts • & diaphragm is raised. • This decreasesvolume • of the lungs. • Decreasing volume • causes pressureto • increase. • Air is expelled out of • lungs to equalize • pressure.

37. Boyle's Law applied • Explains pressure differences that drives breathing • How hypodermic syringes & straws can be filled • Why soda pop cans fizz when you open them • “The Bends” in scuba diving Divers breath compressed air

38. Chap 7, Clicker Question #3: Determine what the volume will be in the lungs of a scuba diver if: • Initial pressure is at 3 atm(66 feet down) • sea level P is = 1atm • Vol of air in lungs is initially 4 liters • Assume Temp is constant As diver races to the surface rapidly, & holding his breath, what happens to the volume of his lungs?

39. Chap 7, Clicker Question #3: Determine volume of lungs: • Initial pressure is at 3 atm(66 feet down) • sea level P is = 1atm • Vol of lungs is initially 4 liters • Temp is constant • 0.75 L • 2 L • 12 L • 44.1 L

40. Use Boyle’s law: P1V1 = P2V2 Rearrange to solve for V2 V2 = P1V1 = (3 atm) (4 liters) = 12 liters P2 1 atm The lungs will burst (6 Liters maximum) and cause an embolism! Now what about temperature in all of this? We have to wait for another invention!

41. Galileo (1592): 1st thermometer based on expansion/contraction of air.  Fahrenheit (1714): 1st mercury thermometer FP of water (32°) to body temp (~100°).  Celsius (1742) centigrade scale BP water (0°) to the FP of water (100°) [inverted eventually] Now: BP of H2O = 100°C FP of H2O = 0°C

42. Lord Kelvin extrapolates to determine Absolute Zero Gay-Lussac’s graph lines extrapolated predict gasses would have zero volume at a temp of -273.15 °C (all gases liquefy or solidify before this low temperature is reached)

43. William Thomson, Lord Kelvin, (1848) used Celsius’ degree size but started at zero Kelvin (-273°C).

44. Conversions from one scale to the next: oC to oF muliply by 1.8 and add 32 e.g. 100 oC x 1.8 oF = 180 oF 1oC 180 oF + 32 = 212 oF oF to oC subtract 32 and divide by 1.8 e.g. 72 oF - 32 = 40 oF 40 oF x 1 oC = 22 oC 1.8 oF

45. Conversions from one scale to the next: oC to K just add 273 e.g. 100 oC + 273 = 373 K K to oC subtract 273 e.g. 293 K – 273 = 20 oC 20 oC x 1.8 oF = 36 oF 1oC 36 oF + 32 = 68 oF

46. Charles's LawTemperature-Volume: Direct Relationship V = constant x T or V / T = constant Pioneer balloonist Jacques Charles equal volumes of all gases expand equally with the same increase in temperature

47. Volumeof a fixed amountof gas maintained at constant pressureisdirectly proportional to its absolute temperature

48. 7.3B Charles’s Law:How Vol & Temp of a Gas Are Related • If one quantity increases, the other increases • Dividing vol by temp is a constant, k. V Volume = k = constant T Temperature Temperature must be expressed in Kelvins.

49. Charles’s Law If temp of gas in cylinder is doubled, volume of gas inside cylinder doubles. • This behavior can be explained by the equation: V1 V2 = T1 T2 initial conditions new conditions

50. Sample Problem 7.3