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Chapter 6: The States of Matter

Chapter 6: The States of Matter. Suggested Problems :. States of Matter. ______ Definite volume and shape ______ Definite volume but not shape Takes the shape of its container _______ No definite volume or shape Will not only take the shape of its container it will fill it completely.

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Chapter 6: The States of Matter

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  1. Chapter 6: The States of Matter Suggested Problems:

  2. States of Matter • ______ • Definite volume and shape • ______ • Definite volume but not shape • Takes the shape of its container • _______ • No definite volume or shape • Will not only take the shape of its container it will fill it completely

  3. Solids • A __________ state of matter • Atoms (or molecules) are “touching” • Strongest intermolecular forces • Hold atoms (or molecules or ions, etc.) rigidly in a 3D crystalline lattice

  4. Liquids • Also a __________ state of matter • Atoms (or molecules) are “touching” • Intermolecular forces hold atoms (or molecules) in contact, but not rigidly in place…molecules can slide past each other

  5. Gases • Virtually ___ intermolecular forces • Gaseous molecules (etc.) comprise a very small percent of the sample volume • Gaseous molecules are in constant random motion…the velocity is related to temperature • Molecules collide with walls of container and with each other and bounce off with no loss of energy

  6. Properties of Gases • Gases are the best understood state of matter, because we ignore intermolecular forces • The volume a gas sample occupies is a function of three variables: • ___________ • ___________ • ___________

  7. Pressure • Pressure is force per unit area • Units of pressure: • Pounds per square inch • Torr or mmHg • Atmosphere 1 atm = 760 torr = 760 mmHg = 14.7 psi

  8. Measuring Pressure: Barometer gravity pulling down air pushing down

  9. Pressure Conversion Example • The gauge on an oxygen gas cylinder reads 1272 psi. Express this in atm and torr. (1 atm=14.7 psi)

  10. Volume of a Gas • Imagine a fixed amount of air at a given temperature and pressure in a balloon • What will happen to the volume if we add more air?

  11. Volume of a Gas • Imagine a fixed amount of air at a given temperature and pressure in a balloon • What will happen to the volume if we squeeze the balloon (increase pressure)?

  12. Volume of a Gas • Imagine a fixed amount of air at a given temperature and pressure in a balloon • What will happen to the volume if we increase the temperature?

  13. The Combined Gas Law • P is pressure • V is volume • T is ________ temperature • n is number of moles

  14. The Empirical Gas Laws • Boyle’s Law • Volume is inversely proportional to pressure (constant n and T) • Charles’s Law • Volume is directly proportional to the Kelvin temperature (constant n and P) • Avogadro’s Law • The volume of a gas is directly proportional to the number of gas moles (constant T and P)

  15. Boyle’s Law: Example • 15 liters of argon is collected at an initial pressure of 1.05 atm. If the gas is compressed to a new pressure of 1510 torr, what is the new volume?

  16. Charles’s Law: Example • The temperature of 35.6 mL of neon is increased from –35.4ºC to 75.2ºC. What is the new volume?

  17. Combining Boyle’s and Charles's Laws • A bubble of air having a volume of 75.0 mL is released from 35 feet under the sea (where the pressure is 2.07 atm and the temperature is 18 ºC). What will the new volume be at the surface, where P=0.967 atm and T=23 ºC?

  18. The Ideal Gas Law • Combines the elements of the ________ gas laws

  19. Standard Conditions • STP = standard temperature and pressure T = _____ K (_____ ºC) P = _____ atm = ____ torr

  20. Example • What volume will 2.0 grams of helium occupy at a temperature of 290K and a pressure of 800 torr?

  21. Which Gas Law to Use? • Use the combined gas law when the problem describes two sets of conditions • the pressure and/or temperature changes • Use the ideal gas law when there are a single set of conditions

  22. Nitrogen 0.75 atm Oxygen 0.25 atm Dalton’s Law of Partial Pressures • The ______ pressure of a gaseous mixture is the sum of the partial pressures Empty container Nitrogen 0.75 atm Oxygen 0.25 atm Ptotal=

  23. Graham’s Law • __________ is net movement of a gas from an area of high concentration (pressure) to an area of lower concentration • __________ is the movement of a gas through a pinhole • Both Diffusion and Effusion follow Graham’s Law Rate is an amount per time

  24. Graham’s Law Example • Oxygen Molecules weigh 16 times as much as hydrogen molecules. Which molecule will diffuse faster and how much faster?

  25. Changes of State Gas condense vaporize sublime condense freeze Liquid Solid melt What do all of these changes in state have in common?

  26. Energy • Energy is the ability to do work • Kinetic Energy: energy due to motion • _________ Energy: stored energy • Heat Energy: the sum of the kinetic and potential energies of molecules in a sample

  27. Energy and Its Units • calorie (cal): is the amount of heat needed to raise the temperature of 1 gram of water by 1 degree Celsius at 15 oC • kilocalorie = 1000 cal • A food Calorie = 1000 cal • Joules (J): are the metric unit of energy 1 cal = 4.184 J

  28. Energy Conversion Example • A candy bar has 350 Calories. How many joules does one candy bar contain?

  29. Heat and Temperature • __________: is a measurement of the average kinetic energy of the molecules in a sample • ___ is measured in degrees with a thermometer • _______: is the sum of the kinetic (and potential) energies in a sample • _____ is measured in calories with a calorimeter

  30. Calorimeters Bomb Calorimeter Coffee Cup Calorimeter

  31. Specific Heat • Specific heat (SH) is the amount of heat needed to raise the temperature of 1 gram of material by one degree Celsius

  32. Specific Heat Example • A 10.0 gram sample of copper at 25 ºC has a final temperature of 100 ºC when 289 J of heat are added. What is the specific heat of copper? (SH of liq water = 4.18 J/goC)

  33. Average KE Distribution of Energy • In a sample of material, the kinetic energies of the molecules follow a Boltzman Distribution: # Molecules Kinetic Energy

  34. Kinetic Energy Distribution Thigh Tlow # Molecules Kinetic Energy

  35. Changes of State Gas condense vaporize sublime condense freeze Liquid Solid melt

  36. Vaporization • The most energetic molecules in a liquid have sufficient kinetic energy to escape into the ____ phase • Once the molecules are free as gases, they exert a pressure • This is called the ______ pressure • How does vapor pressure depend on temperature?

  37. Vapor Pressure of Water and Ethanol

  38. Boiling Point • The boiling point of a liquid is the temperature where the vapor pressure equals the ambient pressure. • The _______ boiling point of a liquid is the temperature where the vapor pressure equals 760 torr. • How does boiling point depend on pressure?

  39. Changes of State Gas condense vaporize sublime condense freeze Liquid Solid melt What do all of these changes in state have in common?

  40. Freezing/Melting Point • The melting point of a substance is the temperature at which a crystalline solid changes to a liquid. • What is the difference between melting point and freezing point?

  41. Energy Changes and Changes of State • Imagine recording the temperature of an 18 gram (i.e., 1.0 mole) sample of ice at -40ºC as heat is added No T No T

  42. Heating Curve for 1 Mole of Water Water is boiling: Heat of vaporization 40.7 kJ/mol Ice is melting: Heat of fusion 6.02 kJ/mol

  43. Molar Heat of Fusion • DHºfusis the heat required to convert one mole of solid to a liquid at at its normal melting point • DHºfus represents the energy needed to break down intermolecular forces and allow molecules to slide around the liquid phase

  44. Molar Heat of Vaporization • DH°vap is the heat required to convert one mole of liquid to a gas at at its normal boiling point • DH°vap represents the energy needed to break intermolecular forces and allow molecules to escape into the gas phase

  45. Putting it all Together • How much heat is required to convert an 18 gram piece of ice at -40 oC to steam at 120 oC?

  46. Heating Curve for 1 Mole of Water D E Heat of vaporization 40.7 kJ/mol C q = m(SH)(T) SH ice = 2.1 J/goC SH liq = 4.18 J/goC SH gas = 2.0 J/goC B Heat of fusion 6.02 kJ/mol A

  47. Question • Explain why orange growers spray their trees with water when there is a threat of freezing temperatures.

  48. Question • Why does steam at 100ºC cause more severe burns than water at the same temperature?

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