Gas Pressure

1. Gas Pressure Pressure = Force/Area Eg: lbs/in2, kg/m2, g/cm2 Force applied to a reduced area will INCREASE Pressure Increased Force applied to same Area will INCREASE Pressure

2. Measuring Pressure • Atmospheric pressure: • Gases have MASS..that is pressing down on you. • High Pressure compresses air leading to better weather. • Low Pressure allows for condensation. RAIN..Clouds. Barometer: Measures atmospheric pressure using mercury. 1 atm = 760 mm Hg. = 760 torr. (Evangelista Torricelli) = 14.7 lb/square in. ( English Units)

3. Converting Units of Pressure • Convert to atmospheres • 385 mmHg • 570 torr

4. More Conversions • What is equivalent of 0.930 atm in mm mercury and in torr.?

5. Measuring Diff. Pressure • Manometer: used to measure the pressure of gas in a closed container. • Differential Manometer; measures pressure by revealing the difference between two pressures. • Gas Gauge: Membrane separates gas and a springlike mechanical device exerts a force opposing the gas . This force is calibrated to read Pounds Per Squre inch and mm. Hg

6. The Gas Laws • The Physical Properties of gases depends on 4 Variables: • Pressure (P) • Volume (V) • Temperature (T) • Moles (n)

7. The Laws • Varying 2 properties and keeping 2 constant allows us to observe many behaviours of gases. • Boyle’s Law: Vary pressure and volume • Charle’s Law: Temperature and Volume • Gay Lussac’s Law: Temperature and Pressure • Avogadro’s Law: moles and Volume

8. Boyles Law • The product of a gas pressure multiplied by it’s volume is CONSTANT. • Experimentation reveals: • As pressure decreases, the volume increases, or as the pressure increases, the volume decreases. • The P x V product is Constant. • Therefore: • P1V1 = P2V2

9. Breathing • As lung volume increases, gas pressure in lungs decreases, and air will flow into the lungs. • As lung volume decreases, gas pressure increase and air will flow out of the lungs.

10. Examples • A 712 mL sample of gas at 505 torr is compressed at constant temperature until its final pressure is 825 torr. What is it’s final volume.

11. Another example • A 2.0 L sample of gas at 0.800 atm must be compressed to 1.6L at constant temperature. What pressure in atmospheres must be exerted to bring it to that volume?

12. Charles Law • V1/T1 = V2/T2 A 512mL sample of a gas, in a cylinder with a moveable piston at 0 deg. C is heated at a constant pressure of 0.800 atm to 41 deg. C. What is its final volume?

13. Gay Lussac’s Law • As temperature increases, pressure increases • The ration P/T is a constant, the the ration P/t is not constant. (t is in deg. C, T is in Kelvin) • Kelvin = Celsius + 273 • P1/T1 = constant = P2/T2

14. Example • A constant volume sample of gas at 27 deg. C and 1.00 atm. Is heates so that its pressure increases to a final value of 1.5 atm. What is its final temperature?

15. Avogadros’s Law • V = constant = n • Volume of gases at the same temperature and pressure contain equal numbers of molecules (n) • V1/n1 = V2/n2

16. Combined Gas Law • P1V1/T1 = P2V2/T2 What is the new pressure of a gas when both temperature and volume are changed simultaneously? What is the new T of a gas when P and V are changed? What is the new V of a gas when T and P are changed?

17. Ideal Gas Law • PV / nT = constant • PV = nRT • R is the universal gas constant. • At 1 atm, 273K, 1.0 mol, 22.4 L • R = 0.0821 Latm/Kmol

18. What gas is better? • To fill your tires with? • To fill your cushions?

19. Ideal Gas Law and Molar Mass • n = g/M ; • Number of moles (n) = sample in grams/Molar Mass • (g)/ (g/mol) = mol • M = gRT/PV

20. Dalton’s Law of Partial Pressures • Ptotal = Pa + Pb Air Pressure = 760 mm Hg or Torr, Partial Pressure of Oxygen is 160 Torr, What is Nitrogen?

21. Henry’s Law • The amt. of gas that will dissolve in a liquid depends manly on the gas pressure and is described by Henry’s Law. • As gas pressure increases, more gas will dissolve. • (Volume of gas dissolved)/(volume of liquid) = CH x gas pressure

22. Chapter 6 • Solids: fixed volume and shape • Liquids: Fixed volume but no shape • Gases: volume depends on the the volume of container. A gas will always occupy all of its container.

23. Kinetic Energy at Molecular level • Kinetic Energy: energy of a moving body, or energy of motion. On a molecular level, kinetic energy depends on temperature. • T increases, so does Kinetic energy. • Molecules move faster • Kinetic Energy tends to override attractive forces.

24. Phase Transitions • Melting: solid transforms into a liquid with the addtion of sufficient heat. • Freezing: reverse process • MP and FP of a solid are identical. • Vaporization: transition from liquid to gas. • Condensation: reverse process • Phase transitions are reversible and occur in either direction.

25. Physical Properties • Molar Heat of Fusion: The heat required per mole of solid to melt it . • Molar Heat of Vaporization: The heat per mole of liquid required for vaporization.

26. Attractive Forces between Molecules • Van der Waals Forces, or intermolecular , or secondary forces… all the same. • They cause gases to condense and liquids to freeze into solids. • They are electrical and they arise from molecular polarity.

27. London Forces of Interaction • Even for molecules that have no permanent polarity, temporary dipoles exist for brief moments. • Temporary dipoles are caused by erratic motions of electrons that result in uneven distributions of electric charge . • The attractive force resulting from this temporary dipole is called a London Force.

28. Dipole – Dipole Interactions • Molecules with permanent dipole moments, or POLAR Molecules also attract different molecules possessing dipole moments.

29. Dipole – Dipole Forces • The attractive forces between 2 polar molecules or mixtures of the 2 are called dipole-dipole forces. • Substances composed of polar molecules have higher melting and boiling points then those composed of molecules possessing no dipole moment.

30. The Hydrogen Bond • The difference between Hydrogen and each of 3 elements (F, O2, N2) electronegativity is very large. This polarity leads to a unique strong attractive force called : Hydrogen Bond. • Eg. H20, NH3, HF

31. The Hydrogen Bond (Water)

32. Surface TensionResisting the expansion of the liquid’s surface

33. Surface-Active Agentsaka: Surfactants • Reduce surface tension of water. • Soaps • Detergents • Creation of stable foams. • Consist of both ionic and • Non polar segments

34. Amphipathic MoleculesHydrophilic- Water Loving and Hydrophobic – Water Hating

35. How a Surfactant reduces surface tension. • Surfactants tend to stay at surface of water • At the interface, hydrophobic end can escape the water. • Hydrophilic end remains nestled in water. • The concentration of the water at the surface is reduced. Therefore so are the attractive forces between water molecules. • Surface tension is reduced.

36. Molecular Mixtures • Solution: a mixture of substances with different kinds of molecules (or ions) uniformly distributed visually throughout the mixture. • Secondary forces are key to the formation of solutions. • Like dissolves like (similar secondary forces)

37. Molecule A & B in Solution

38. Predicting Solutions • Can water and ethanol go into solution? • 1. Draw Structural formula • 2. Are they polar or nonpolar or both? • 3. They are both polar. • 4. Do they have any Hydrogen Bonds? Yes OH group. • 5. Secondary forces are about the same, they will easily bond with each others secondary forces. Therefore …YES it will go into soluton!

39. More Examples • Carbon Tetrachloride and Water? • Weak 2ndary forces will not combine with strong secondary forces. • Hexane and Water? • Hexane is in liquid state from London forces • Water is Hydrogen bonding. • NO…solution

40. The Vaporization of Liquids • Vapor: the gaseous part of a liquid and gas simultaneously present. • Gas in a gas station ( the smell is a gas vapor) • Oxygen vapor exists at -200 deg. C • To vaporize the molecule must free itself of neighbor’s attractive forces and enter the gas phase.

41. Influence of Secondary Forces on Vapor Pressure • VP measures the escaping tendency of a substance’s molecules. • The smaller the forces, the greater tendency of molecules to escape & the greater the vapor pressure. • The larger the forces, the smaller the vapor pressure.

42. Boiling • Bubbles of vapor form throughout the liquid in an open container. • A liquids molecules must overcome the opposing force of the atmospheric pressure to enter the gas space over the liquid. • When the temperature of the liquid is such that it’s vapor pressure is less than atmospheric pressure, vapor will leave only at its’ surface. • When VP is equal to atmospheric pressure, bubbles of vapor form throughout the bulk of the liquid. BOILING • It’s all about VaporPressure of Liquid and the Atmospheric pressure being equal….it changes with atomospheric pressue.

43. Normal Boiling Point • Temperature at which boiling occures under an external pressure of exactly 1 atm. • We can use normal boiling point as a good qualitative indicator of the secondary forces in a liquid.

44. Using Molecular structure to predict boiling points • H2, CH4, NH3, HF • 1) Are there any polar compounds? • Polar compounds have higher boiling points then non-polar compounds. • Why? Non-polar compounds are only influenced by London Forces. • 2. Do any of the Polar Compounds have H-bonds? The one with the highest electronegativity will have the highest boiling point. • 3. If they have similar forces and electronegativities, then which molecule is bigger/heavier, and more electrons and therefore will interact with stronger London forces then a smaller molecule.

45. Stronger Bonds, Higher Boiling Point • What intramolecular force is stronger? • Ionic Bond, or Covalent Bond? • Covalent Bonds are stronger. • Ionic bonds will dissassociate into ions in water. They are not actual molecules. • What is stronger NaCl or Graphite (long non-polar carbon chain?)

46. Equilibrium In a closed system, the rate of vaporization = rate of condensation Vapor pressure in equilibrium is called Equilibrium Vapor Pressure. Dynamic Equilibrium: Liquid + Heat -----------vapor vaporization Vapor --------------liquid + heat condensation

47. Evaporation • Vaporization in an Open System. • There will be an imbalance with no condensation, so the system is not in equilibrium. • Therefore the liquid must continuously absorb heat from the surroundings. If the flow of heat is restricted, the temp. of the liquid will drop. • Evaporation is a cooling process.

48. Using Secondary forces to predict boiling points • H20(water) vs H2S (hydrogen sulfide) • Boiling point of water is 100 deg. C • Boiling point of H2S is -61 deg. C • H bonds only happen with O, N, and F. • H2S is limited to interact with weaker secondary forces.

49. example • Which has a higher boiling point? • Ammonia (NH3) or Phosphine (PH3)

50. Attractive Forces and The Structure of Solids • Crystalline Solids: • Solids form crystals arranged in regular patterns. They can exist in well-defined shapes, like prisms, cubes… • They have sharp edges or flat faces because it is that highly ordered inner surface that is exposed. • Amorphous Solids: • Lacking crystalline structure. • Eg. Non crystalline form of SiO2(glass)