Chapter 5 Gases

# Chapter 5 Gases

## Chapter 5 Gases

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##### Presentation Transcript

1. Chapter 5 Gases

2. The nature of gases Indefinite shape and indefinite volume expand to fill their containers Compressible Fluid – they flow Low density 1/1000 the density of the equivalent liquid or solid Properties of a gas stem from atoms/molecules being far away from each other, great amount of space in between

3. Kinetic molecular theory • Kinetic molecular theory explains the gas laws • A gas is a collection of particles (atoms or molecules) in constant motion • A single particle moves in a straight line till it collides with the wall or another particle

4. 3 postulates of KMT • 1. The size of a particle is negligibly small • Since the space between atoms is very large in comparison to the atoms themselves • 2. The average kinetic energy of a particle is proportional to the temperature in kelvins • Heat/Thermal energy causes movement of the molecules • The higher the temperature, the higher the average kinetic energy • 3. The collision of one particle with another (including container walls) is completely elastic • There is no loss of energy in elastic collisions • Energy can be transferred, but overall there is no net loss

5. Pressure • Is equal to force/unit area • P=F/A • Is caused by the collisions of molecules with the walls of a container

6. Vacuum pumps • Vacuum pumps – can remove gas from a container, in turn reducing the pressure of that container • Useful for performing air free chemistry (mainly to avoid H2O and O2 in the air)

7. GAS LAWS Boyle’s Law Charles’ Law Gay-Lussac’s Law Avogadro’sLaw Ideal Gas Law Daltons Law

8. Gas law intro • Each law details the relationship between the following properties of gases • Volume • Pressure • Temperature • Moles • Each law is an equation • You must memorize all of the gas law equations • Solving problems typically involves rearranging the base equation to solve for a specific unknown

9. BOYLE’S LAW The relationship between pressure and volume

10. Boyle’s law • Robert Boyle – 1627-1691 • Boyle’s Law - The volume of a gas is inversely proportional to the pressure applied to the gas when the temperature and moles are kept constant. • Decrease in volume = Increase in pressure. • Increase in volume = Decrease the pressure

11. The relationship between temperature and volume

12. Charles’ law • Jaques Charles – 1746-1823 • At a fixed pressureand moles, the volume of a gas is directly proportional to the temperature of the gas • As the temperature increases, the volume increases • As the temperature decreases, the volume decreases • TEMPERATURES MUST BE IN KELVIN

13. Gay-Lussac’s law The relationship between pressure and temperature

14. Gay-Lussac’s Law • Joesph Louis Gay-Lussac – 1778-1850 • At a fixed volumeand moles, the pressure of a gas is directly proportional to the temperature of the gas • As the temperature increases, the pressure also increases • As the temperature decreases, the pressure also decreases • TEMPERATURES MUST BE IN KELVIN

15. Avogadro’s lawThe relationship between volume and moles

16. Avogadro’s law • AmedeoAvaogadro – 1776-1856 • At a fixed pressureand temperature, the moles of a gas is proportional to the volume of the gas • As the molesincreases, volume also increases • As the molesdecreases, volume also decreases • n = moles

17. The Ideal Gas Law

18. Ideal gas law • Combining all previous relationships into one equation • The ideal gas law is most often written as • R - universal gas constant • R = 0.082 L x atm x K-1 x mol-1 • R = 62.36 L x torr x K-1 x mol-1 • R = 62.36 L x mmHg x K-1 x mol-1 • R = 8.315 L x kPa x K-1 x mol-1 • ALL TEMPERATURES MUST BE IN KELVIN!!

19. Molar volume and STP • One mole of any gas occupies exactly 22.4 liters (dm3) at STP. • STP = Standard Temperature and Pressure • Temp = 0ºC = 273K • Pressure = 1atm = 760 mmHg = 760 torr etc. • This is often referred to as the “molar volume” of a gas.

20. Equal volumes of gases, at the same temperature and pressure, contain the same number of particles, or molecules • Thus, the number of molecules in a specific volume of gas is independent of the size or mass of the gas molecules • Therefore, the rules applies the same to all gases

21. Problems • A cylinder contains 5.00L of gas at 225K. If the temperature is increased to 345K, what will the new volume be? • Rosy has a 5.00L tank of H2 gas. If the pressure inside the tank is 800.0torr and the temperature is 300.0K, how many moles of hydrogen does her tank contain?

22. Density of a gas • Density = mass/ volume • What is the mass of a gas? • M = molar mass of gas (grams/mole) • m = mass of gas (grams) • n = moles of gas (moles) • V = volume of gas (Liters) • dgas = density of the gas

23. Density of a gas • Make sure to choose the R value that cancels out the correct pressure units • g/L should be the final density units

24. Dalton’s LawThe Law of Partial Pressures

25. Dalton’s law • John Dalton – 1766 - 1844 • Dalton’s Law: The total pressure exerted by a mixture of gases is the sum of the individual pressures of each gas in the mixture • Dalton’s Equation:

26. Deep-sea diving and partial pressures • Lungs need to breathe oxygen within a certain partial pressure range 0.21 -1.6 atm • Hypoxia – O2 partial pressure to low, not enough O2 • Oxygen toxicity – O2 partial pressure too high, too much O2

27. Problems • The air in this room contains the gases shown below at their respective partial pressures in kPa. What is the pressure your body feels from the air in this room?

28. Mole fraction (χa) • In a gas mixture, there is a certain amount of moles of each individual gas • The number of moles of one of those gases divided by the total number of moles from all the gases is the mole fraction • χa = mole fraction • na= moles of a • ntotal = total moles of gas in gas mixture

29. Problem • If you had 101.3 moles of air, what is the mole fraction of O2 in that sample?

30. Stoichiometry with gases • General stoichiometry • Solution Stoichiometry • Gas Stoichiometry

31. Molar volume and stoichiometry • 1 mole of gas at STP = 22.4 L • An easy way to get to moles at STP • In general • Gas reactions still apply the same stoichiometry concepts • The chemical equation relates the reactants and products • To use the chemical equation mole ratios you need to convert reactants/products to moles • This lets you go from moles of one compound to another • The only new concept is that the ideal gas law allows us another way to get to moles

32. Problem CO(g) +2H2(g)  CH3OH(g) What volume in liters of hydrogen gas, at 355K and a pressure of 738 mmHg, is needed to synthesize 35.7g of methanol (CH3OH)?

33. Temperature and Molecular Velocities • Gas particles at a given temperature have the same average kinetic energy • Lighter particles (Hydrogen) have faster average velocities than heavier particles (Argon) to compensate for lack of mass, so the kinetic energy can be equal

34. Root mean square velocity (urms) • urms– the root of the average of the squares of the particles velocities • R = universal gas constant = 8.314 (J/(mol*K)) • M = molar mass of gas IN KILOGRAMS • O2 molar mass = 32g/mol = 0.032kg/mol • J = (kg*m2)/(s2), now units of kg match • T = temperature in Kelvin

35. Diffusion • Diffusion- describes the mixing of gases. The rate of diffusion is the rate of gas mixing • Lighter gas particles diffuse faster than heavier gas particles

36. Effusion • Effusion- describes the passage of gas into an evacuated chamber • Lighter particles effuse faster than heavier particles • Grahams law of effusion – rate of effusion of two gases is related

37. Real gases • Ideal gas laws are suited for conditions close to STP and are based off of the 3 kinetic theory postulates • Negligible particle size, average KE related to T, elastic collision • At lower temperatures and higher pressures that Ideal gas laws begin to become less effective

38. Real gas - Volume • Particle size has a greater impact at higher pressures (lower volumes) • This causes the actual volume to be greater than predicted with the ideal gas law

39. Correction for real behavior • n = moles • b = constant dependant on gas, use chart 

40. Real gas - Pressure • Lower temperatures allow for intermolecular forces to have a greater effect, decreasing the collisions that occur • This causes the actual pressureto be less than predicted with the ideal gas law

41. Correction and van der Waals equation • Combining the pressure and volume corrections give the van der Waals equation

42. Chapter 5….. out of gas.