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  • 1. Chem Lab 1 Water
  • 2. Lab 1.1 States of Matter
    • Observations and Inferences
      • Solids
      • Liquids
      • Gases
    • Phase Changes
  • 3. SOLIDS
    • Particle Spacing
    • Observation
      • The solid cannot be compressed
    • Inference
      • The particles must be as close as they can get.
    • Freedom of Movement
    • Observation
      • Solids maintain their shape
    • Inference
      • The particles must be held in fixed positions.
  • 4.  
  • 5. LIQUIDS
    • Particle Spacing
    • Observation
      • The liquid cannot be compressed
    • Inference
      • The particles must be as close as they can get.
    • Freedom of Movement
    • Observation
      • Liquids take the shape of their container.
    • Inference
      • The particles must be free to move to new positions.
  • 6. Ice- water molecules in fixed positions Water molecules free to move
  • 7. GASES
    • Particle Spacing
    • Observation
      • The gas can be compressed to a smaller volume
    • Inference
      • The particles must be widely spaced.
    • Freedom of Movement
    • Observation
      • Gases take the shape of their container.
    • Inference
      • The particles must be free to move to new positions.
  • 8. Gas particles widely spaced Liquid particles as close as can get but free to move
  • 9.  
  • 10. BREAKING AND FORMING ATTRACTIONS
    • Forming attractions
      • Releases energy
      • Compare to letting go of a stretched spring
      • Energy is released as PE is converted to KE
      • as spring comes back together
    • Breaking attractions
      • Requires energy
      • Compare to pulling apart a spring
      • Energy must be transferred in to do work to change the position of the spring increasing its potential energy
  • 11.  
  • 12. Attractions form releasing energy
  • 13. Phase Changes
    • Solid to Liquid
      • What change in particles is required?
        • Particles must be broken out of fixed positions. They gain freedom of movement.
      • Are attractions broken or formed?
        • Attractions must be broken.
      • Is energy transferred in or out of the system?
        • Energy must be transferred into the system.
  • 14.  
  • 15. Melting
    • Is this change endothermic or exothermic?
      • Since energy must be transferred in, the change is endothermic.
    solid liquid enthalpy net change in enthalpy
  • 16.  
  • 17. Boiling
    • Energy must be transferred into the system to:
      • break the attractions between molecules
      • spread the molecules apart.
      • The widely spaced molecules have more potential energy.
    • Is this change endothermic or exothermic?
      • Since energy must be transferred in, the change is endothermic.
  • 18. Liquid to Gas solid liquid enthalpy net change in enthalpy
  • 19. Enthalpy
    • The change in enthalpy is measured as the energy transfer between the system and surroundings
    • Enthalpy is related to the potential energy of the reactants and products
  • 20. POSITIVE CHANGE IN ENTHALPY
    • A positive change in enthalpy
      • Indicates an increase in PE
      • Products
        • have more PE than the reactants
        • and are less stable than the reactants
      • Energy must be transferred into the system from the surroundings to increase the PE
      • ENDOTHERMIC reactions are NOT favored based upon enthalpy
  • 21.  
  • 22. NEGATIVE CHANGE IN ENTHALPY
    • A negative change in enthalpy
      • Indicates an decrease in PE
      • Products
        • have less PE than the reactants
        • and are more stable than the reactants
      • Energy must be transferred out of the system to the surroundings to decrease the PE
      • EXOTHERMIC reactions are favored based upon enthalpy
  • 23.  
  • 24.  
  • 25. ENTROPY
    • Associated with the disorder
    • Liquids are more disordered than solids
    • Gases are more disordered than liquids
    • A greater freedom of movement leads to more possible random locations- a greater disorder
    • An increase in entropy or disorder is favored.
  • 26. Entropy There are two natural tendencies behind spontaneous processes: the tendency to achieve a lower energy state and the tendency toward a more disordered state
  • 27. The greater the number of configurations of the microscopic particles (atoms, ions, molecules) among the energy levels in a particular state of a system, the greater the entropy of the system
  • 28. Water Molecule
    • 2 hydrogen atoms and 1 oxygen atom are held together by covalent bonds
    • MOLECULES
      • are held together by covalent bonds
    • COVALENT BOND
      • Two atoms share a pair or pairs of electrons
  • 29.  
  • 30. ELECTRICAL FORCE
    • Force of attraction and repulsion due to charge
      • Directly related to charge
      • Inversely related to distance between charges
    • Opposite charges are attracted
    • Like charges repel
    • Negative electrons are attracted to positive protons within the nucleus
  • 31. Oxygen atom has 8 positive protons pulls harder on neg e- e- spend more time This end becomes partially negative Hydrogen atom has only one proton less pull on neg e- e- spend less time This end becomes partially positive δ + δ -
  • 32. Polar Covalent Bond
    • Covalent bond
      • Sharing pair(s) of electron
    • Unequal sharing
      • due to differences in the atoms’ pull on electrons
      • lead to oppositely charged ends (poles)
  • 33.  
  • 34. Electron Density in Water Molecule Notice the electron density is greater around the oxygen. The electrons are more likely to be found around the oxygen atom than the hydrogen atom
  • 35. Dipole-Dipole Attraction
    • Attraction between polar molecules
    • dipole - oppositely charged ends
    • Negative oxygen-end of one water molecule attracted to positive hydrogen-end of another water molecule
  • 36. Dipole-dipole attractions between water molecules
  • 37. Notice the polar covalent bonds are shorter (and thus stronger) than the dipole-dipole attractions.
  • 38.  
  • 39. Surface Tension in Water is the result of dipole-dipole attractions That’s why the bug can walk on top of the water. The water molecules are attracted to each other
  • 40. Surface Tension at the particle level
  • 41. Water drops are high and round due to the attractions between the water molecules
  • 42. Lab 1.2 Measuring the Heat Transfer to Melt Ice
  • 43. ICE- open crystal lattice structure
  • 44. Measuring Heat Transfer
    • Use a calorimeter
      • Styrofoam cup
      • Thermometer
      • Water
    • Calculate the heat transfer
      • q = mC  T
      • m = Mass of water in grams
      • C = Specific heat (water 4.18 J/gK)
      •  T = Change in Temperature
  • 45. Procedures
    • Measured the mass of water in calorimeter
      • Measured mass of empty cup
      • Measured mass of cup with hot water
    • Measured change temperature
      • Measured initial temp of hot water in cup
      • Measured final temp after ice melts
    • Measured the mass of ice
      • Measured mass of cup with melted ice and original hot water
  • 46. hot water thermometer ice Calorimeter to measure heat transfer
  • 47.  
  • 48. Calorimetry Calorimetry is a technique used to measure heat exchange in chemical reactions A calorimeter is the device used to make heat measurements EOS Calorimetry is based on the law of conservation of energy
  • 49. Results 159.805 g Mass of cup with melted ice and original hot water 19.2 °C Final temperature of hot water 68.6 °C Initial temperature of hot water 99.722 g Mass of cup with hot water 3.867 g Mass of empty cup
  • 50. Precision & Accuracy Illustrated
  • 51. Significant Digits
    • All digits in a number that are known with certainty plus the first uncertain digit
    • The more significant digits obtained, the better the precision of a measurement
    • The concept of significant figures applies only to measurements
    • Exact values have an unlimited number of significant figures
  • 52. Rules for Zeros in Significant Figures Zeros between two other significant digits ARE significant e.g., 1 00 23 A zero preceding a decimal point is not significant e.g., 0.1 00 23 EOS Zeros between the decimal point and the first nonzero digit are not significant e.g., 0.001 00 23
  • 53. Rules for Zeros in Significant Figures Zeros at the end of a number are significant if they are to the right of the decimal point e.g., 0.1 00 23 00 1 0 23. 00 EOS Zeros at the end of a number may or may not be significant if the number is written without a decimal point e.g., 1 000 . compared to 1000
  • 54. Rules for Significant Figures in Calculations KEY POINT: A calculated quantity can be no more precise than the least precise data used in the calculation … and the reported result should reflect this fact EOS Analogy: a chain is only as strong as its weakest link
  • 55. Significant Figures in Calculations EOS 0.762 has 3 sigfigs so the reported answer is 1.39 m 2
  • 56. Significant Figures in Calculations Addition and Subtraction: the reported results should have the same number of decimal places as the number with the fewest decimal places EOS NOTE - Be cautious of round-off errors in multi-step problems. Wait until calculating the final answer before rounding.
  • 57. Calculations
    • Mass of Water in calorimeter
      • cup with water – empty cup
      • 99.722 g - 3.867 g = 95.855 g
    • Change in Temperature of water in calorimeter
      • final temp – initial temp
      • 68.6 °C - 1 9.2 °C = 49.4 °C
    • Mass of ice that melted
      • cup with melt and original water – cup w/ water
      • 159.805 g - 99.722 g = 60.083 g
  • 58. Additional Calculations
    • Heat transfer to melt all the ice
      • q = m C  T
      • 95.855 g • 4.18 J/g °C • 49.4°C
      • = 1.98 x 10 4 Joules
    • Heat transfer to melt one gram of ice
      • heat to melt all ice/mass of ice
      • 1.98 x 10 4 Joules/ 60.083 g
      • = 330 J/g
  • 59.  
  • 60. Endothermic Reaction as Ice Melts
    • System- the ice
    • Surroundings- the hot water in calorimeter
    • Temperature of surroundings decreases
      • Definition of temperature- average KE
      • Since lower temperature
        • The hot water has less KE
        • KE was transferred into the ice- system
  • 61.  
  • 62. Temperature and Phase Change Melting of Ice No change in temp Energy in to break attractions 334J/g Temp (°C) Heat transferred (J) ↑ Temp ↑ avg KE Boiling No change in temp Energy in to break attractions 2,257 J/g
  • 63. Y
  • 64. Entropy Changes associated with the phase changes and temperature changes
  • 65. Lab 1.3 Liquid/Gas Phase Change
  • 66.  
  • 67. In the Flask over the Flame
    • The temperature (average KE) of the liquid water increases
    • Until the boiling point is reached
    • H 2 O (l)  H 2 O (g)
    • Energy transferred into the system from the surroundings (flame)
    • Endothermic reaction
  • 68. H 2 O (l) H 2 O (g) enthalpy BOILING WATER Products are less stable with more PE Reaction not favored based upon enthalpy ENTHALPY
  • 69. Boiling- large bubbles form within the liquid
  • 70. ENTROPY
    • Disorder
    • Entropy increases
    • The gaseous product is more disordered than the liquid reactant
    • Favored based upon entropy
  • 71. In the Flask in the Ice Bath
    • Condensation
    • H 2 O (g)  H 2 O (l)
    • The ice outside the flask melts (Remember the melting of ice requires heat to be transferred into the ice)
    • Energy is transferred from the system (condensing steam) to the surroundings (ice bath)
    • Endothermic reaction
  • 72. H 2 O (l) H 2 O (g) enthalpy CONDENSATION Products are more stable with less PE Reaction is favored based upon enthalpy ENTHALPY
  • 73.  
  • 74. Boiling vs Evaporation
    • Boiling
    • Occurs only at boiling point
    • Large bubbles form at the bottom and move upward
    • Evaporation
    • Occurs at any temperature
    • Occurs only at the surface
    • Molecules on the surface with enough kinetic energy break free of the liquid and move into the gas
  • 75. Temperature- Average KE
    • Boltzmann Curve
  • 76. Another version of Boltzmann curve.
  • 77. Evaporative Cooling
    • Molecules with enough KE can escape
    • Leaving behind molecules with less KE
    • Lowering the average KE of molecules remaining in the liquid
    • Lowering the temperature of the liquid
  • 78. Pressure = Force/area
    • Measured with a barometer
    • Measured in mmHg, psi, N/m
    • Caused by the collision of particles with a unit square
    • Visualize the collisions
    • with the square
  • 79. Collisions of Particles cause Pressure
  • 80. Collisions cause pressure
  • 81. Measuring pressure with a barometer in mm Hg
  • 82. VAPORIZATION- Molecules escape the liquid and move into the gas CONDENSATION- Molecules in the gas collide with liquid and stay in liquid
  • 83. Only water molecules on the surface with enough KE to break dipole-dipole attractions can escape
  • 84. Vapor Pressure
    • Vapor pressure is the result of the collisions of particles that have escaped the liquid and moved into the gas phase
  • 85. Measuring Vapor Pressure Vacuum- no particles in the gas. Liquid vaporizes Proceeds to equilibrium Particles that have escaped from the liquid exert pressure to hold up the column of mercury
  • 86. Vapor Pressure Molecules of the liquid escape the liquid and move into the gas. The molecules collide with the square. This is vapor pressure.
  • 87. Measuring Vapor Pressure
  • 88. Vapor Pressure Equilibrium
    • When a container of liquid is closed, two reversible reactions occur
    • H 2 O (l)  H 2 O (g)
    • Initially, only the forward reaction occurs
    • H 2 O (l)  H 2 O (g)
    • Over time as more molecules escape into the gas, they are more likely to hit the liquid and return
      • Rate of condensation increases H 2 O (g)  H 2 O (l)
      • Rate of vaporization decreases H 2 O (l)  H 2 O (g)
    • Eventually, the rate of the molecules escaping the liquid is equal to rate of molecule returning to the liquid
      • At equilibrium, the rates are equal
      • As system remains at equilibrium
      • conditions remain constant- same vapor pressure
  • 89. Rate Time Vaporization H 2 O(l)  H 2 O(g) Rate decreases Condensation H 2 O(g)  H 2 O(l) Rate increases More particles in gas Equilibrium Rates Equal # molecules in gas stays constant
  • 90.  
  • 91. Vapor Pressure depends upon Temperature
    • As temperature increases
    • The particles have more kinetic energy
    • More particles have enough KE to escape
    • With more particles in the gas,
      • there are more collisions
      • thus vapor pressure increases
      •  Temperature  Vapor Pressure
  • 92. Vapor Pressure and Temperature At a higher temperature, more molecules escape the liquid, thus the vapor pressure is higher
  • 93.  
  • 94.  
  • 95. Boiling Point and Vapor Pressure
    • A liquid will boil when the vapor pressure equals the external pressure (outside air pressure)
    • Bubbles can form within the liquid when the pressure within the bubble (VP) equals the outside pressure of the air
    • If the outside air pressure is changed, water will boil at a temperature other than100 ° C
  • 96. Boiling- large bubbles form within the liquid Bubbles can only form when the vapor pressure of the liquid is equal to outside air pressure Vapor pressure within the bubble keeps the bubble from collapsing
  • 97. In the vacuum, when air particles are removed the outside pressure decreases and water will boil at much lower temperature. The vapor pressure in the bubble will not need to increase as high to equal the outside pressure the water will boil at a lower temperature
  • 98.  
  • 99. Lab 1.4 Decomposition of Water
  • 100. Decomposition of Water Electrolysis The volume ratio of hydrogen to oxygen gas is 2:1
  • 101.  
  • 102. Decomposition of Water Electrolysis
    • 2H 2 O(l)  2H 2 (g) + O 2 (g)
    • O H H O O
    • H H
    • Bonds must be broken (Energy IN)
    • Bonds must be formed (Energy OUT)
    • NET energy transferred INTO system- electrical energy – power supply
  • 103. Decomposition of Water Electrolysis
    • 2H 2 O(l)  2H 2 (g) + O 2 (g)
    • Overall the reaction is endothermic
    H 2 O(l) H 2 (g) + O 2 (g) Net Δ H Net energy transferred IN E IN to break bonds E OUT Bonds form enthalpy
  • 104. Nonpolar Covalent Bonds
    • Hydrogen molecules
      • Two hydrogen atoms held together by covalent bonds
      • Share one pair of electrons
    • Oxygen molecules
      • Two oxygen atoms held together by covalent bonds
      • Share two pairs of electrons
  • 105. Nonpolar Covalent Bonds
    • Same two atoms
    • Exert same pull on shared pair of electrons
    • Electrons spend same time around each atom
    • No oppositely charged ends are formed
    • EQUAL SHARING of e-
      • No dipoles are formed
  • 106. Induced Dipole-Induced Dipole Attractions
    • Attractions between nonpolar molecules
    • With no oppositely charged ends, there is very little attraction between the nonpolar molecules
    • Temporary dipoles are induced as e- density shifts
  • 107. Oxygen molecules O 2 Hydrogen molecules H 2 The reverse reaction- synthesis of water from hydrogen and oxygen gas
  • 108.  
  • 109. Synthesis of Water
    • 2H 2 (g) + O 2 (g)  2H 2 O(l)
    • Overall the reaction is exothermic
    H 2 O(l) H 2 (g) + O 2 (g) Net Δ H Net energy transferred OUT E IN to break bonds E OUT Bonds form enthalpy
  • 110.  
  • 111. Synthesis of Water
    • At room temperature
      • No noticeable reaction occurs
      • When most of the molecules collide they do NOT have enough energy to break bonds
    • With a higher temp- the match
      • Many more of the molecules collide with enough energy to break bonds
      • The reaction is occurs much faster with an explosive release of energy
  • 112.  
  • 113. Just as water flows down the water falls from higher PE to lower PE, so do chemical reactions
  • 114. Collision Model
    • Activation Energy
      • The energy required to break bonds
      • Shown on the uphill portion of the graph
    • Molecules must collide with enough energy to break bonds
      • Effective collisions break bonds
      • A collision without enough energy or not a the right angle will not break any bonds
  • 115. Water Physical Changes Chemical Changes