Chem 102 Part 2

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mct
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212129
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Chem 102 Part 2
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2013-04-10 16:56:21
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After midterm
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Lectures after midterm #13
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  1. Common Ion effect
    Effect of a common ion on solubility
    • Shift in an equilibrium caused by addition of an ion involved in equilibrium
    • -addition of common ion decreases solubility of a sparingly soluble salt
    • -removal of a common ion increases solubility of a sparingly soluble salt
  2. Effect of pH on solubility
    • Removing CO32- (common ion)
    • Increases solubility of CaCO3
  3. Buret reading tips
    • Rinse buret with 5 mL standardized NaOh
    • Buret should be vertical
    • Remove funnel from buret
    • Read bottom of meniscus to + 0.01 mL
    • Eyes should be level with meniscus
    • Buret reading card makes bottom of meniscus darker
  4. Delta Go
    Delta Ho
    Delta So
    • Delta Go = Gibb's free energy
    • < 0 reaction is "spontaneous"
    • > 0 reaction is "nonspontaneous"
    • Delta Go = -RT ln K
    • Delta Ho = Enthalpy
    • < 0 reaction is "exothermic"
    • > 0 reaction is "endothermic"
    • Delta So
    • > 0 increased disorder
    • < 0 decreased disorder
    • Delta Go = Delta Ho - T(delta So)
  5. Energy transferred
    Either as heat and/or work
  6. System
    Part of universe we are focusing on
  7. Surroundings
    Everything else around the system
  8. Internal Energy (E or U)
    • Sum of kinetic energies and potential energies of all particles in a system
    • Typically measure change in internal energy
    • delta E = Efinal - Einitial = Eproducts - Ereactants
    • If system loses energy (delta E < 0), surroundings gain energy
    • If system gains energy (delta E > 0), surroundings lose energy
    • -a delta E = 0, means no energy is being transferred
  9. Delta E = q + w
    • q = thermal energy (heat) - energy transfer that results only in temperature change
    • w = work (mechanical, electrical, etc) = energy transferred when object is moved by force
    • eg. reaction that produces a gas
    • Delta E < 0 energy lost from system to surroundings
    • Delta > 0 energy gained by system from surroundings
    • Delta E = q+w = positive + positive = positive
    • = negative + negative = negative
    • =(positive + negative) = depends
  10. First law of Thermodynamics
    • Law of conservation of energy - total energy of universe is constant
    • Delta E = Delta Esystem + Delta Esurroundings = 0
    • State function - property of system determined by its current state, regardless of the path of that state
  11. State function
    Property of system determined by its current state, regardless of the path to that state
  12. Law of conservation of energy
    Total energy of universe is constant
  13. Enthalpy
    • Consider only work done by expanding gas
    • w = -P(delta V)
    • Most reactions performed at constant pressure:
    • Delta H = Delta E + P(delta V) at const. P;
    • same as qP = delta E - w
    • -change in enthalpy equals the heat gained or lost by system at constant pressure
    • Delta H = delta E reactions not involving gases
    •           reactions with delta nGAS = 0
    • Delta H ~ Delta E reaction with delta nGAS not = 0
  14. Exothermic
    (Delta Hrxn < 0) = heat lost to surroundings
  15. Endothermic
    (Delta Hrxn > 0) = heat gained from surroundings
  16. Delta Hf
    = heat of formation = heat of reaction for producing 1 mole of product from its elements
  17. Delta Hfus
    = heat of fusion for when 1 mole melts
  18. Delta Hvap
    =heat of vaporization for when 1 mole vaporizes
  19. Delta Hdissolution or delta Hsoln
    =heat evoled from dissolving 1 mole of salt into its constituent ions
  20. Calorimetry
    Science of measuring heats of reaction
  21. Calorimeter
    • Device used to exp. determine quantity of heat (q) associated with a chem rxn
    • -"isolated system" where no energy or matter is exchanged with the surroundings
    • -heat of reaction does not escape so must raise temperature of contents of calorimeter
  22. Coffee cup calorimeter
    • Calorimetry performed at constant pressure
    • quantity of heat = specific heat x mass of substance x delta T
    • specific heat x mass of substance = heat capacity
  23. Cold Packs
    • Are based on dissolution of:
    • NH4NO3 (s) -> NH4+ + NO3- aq
    • Delta Hsoln = +25.7 kJ
    • Squeezing cold pack bursts a bag of water, allowing water to mix with NH4NO3 crystals
  24. Self Heating coup cans
    • Based on:
    • CaO (s) + H2O (l) -> Ca(OH)2 (s)
    • Delta H = -65.15 kJ
  25. Hess's Law of Heat Summation
    Enthalpy change of an overall process is sum of enthalpy changes of its individual steps
  26. Standard Heats of Formation
    • Standard enthalpy of formation, delta Hfo = change in enthalpy to form 1 mole of a compound from its elements
    • - characteristic of a substance
    • elements -> 1 mole of compound
    • Standard States - all substances in their standard states
    • Gas - pure gas at 1 atm (~1 bar)
    • Liquid or solid - pure liquids or pure solid
    • Solution - 1 M
    • Element - most stable form (g, l, s) at 1 atm and typically 298K
    • delta Hf (element) = 0
  27. State function
    Property of system determined by its current state, regardless of the path to that state
  28. Second Law of Thermodynamics
    • Entropy (S) - thermodynamic measure of randomness of disorder
    • -as disorder increases, entropy increases
    • -natural progression if from order to disorder
    • Entropy is closely related to probability
    • -spontaneous processes proceed towards states that have highest probability of existing
  29. Second law of thermodynamics definition
    • Processes occur in direction that increases the entropy of the universe
    • delta Suniv = delta Ssystem + delta Ssurroundings
  30. Entropy
    • Thermodynamic measure of randomness or disorder
    • Eg. melting ice - crystalline solid is replaced by less structured liquid
    • vaporization of water,
    • and dissolving NH4NO3 in water
  31. Standard Molar Entropies and Third Law
    • Third law of thermodynamics - perfect crystal has zero entropy at absolute zero (0 K)
    • General trends of entropy in a system:
    • S (solid) < S (liquid) <<< S (gas)
    • -entropy increases when liquids formed from solids
    • -entropy really increases when gases formed from solids or liquids
    • entropy generally increases as:
    • -ngas increases
    • -total volume is increased
    • -temperature is increased
    • -complexity of a molecule
  32. Third law of thermodynamics
    Perfect crystal has zero entropy at absolute zero (0 K)
  33. Entropy Free Energy and Work
    Delta Gosys = delta Hosys - T(delta Sosys)

    • Delta G < 0, reaction is spontaneous
    • Delta G > 0, reaction is non-spontaneous
    • reverse rxn is spontaneous
    • Delta G = 0, reaction is at equilibrium
    • balance between forward & reverse rxns
  34. Coordination Compounds (complexes)
    • [Cu(NH3)6](NO3)2 Coordination compound
    • [Cu(NH3] Comlex ion
    • (NO3)2 Counter Ion
    • list cation, then anion
    • Metal complex goes in [ ]
    • Counter ions outside [ ]
    • Hexaminecopper(II) nitrate
  35. Monodentate ligands
  36. Polydentate ligands
  37. Naming Compounds
    • 1. Name cation first then anion
    • 2. In naming complex, name ligands before metal. List ligands alphabetically, with anionic ligands ending in -o.  Indicate number of ligands using : di-, tri-, tetra-, penta-, hexa- use bis-, tris-, tetrakis-... for ligands that already have the di-, tri- prefixes (ethylenediamine)
    • 3. Count up all anions and look at overall charge to determine oxidation number of metal
    • 4. Designate metal oxidation number using Roman numerals (I, II, III, IV...)
    • 5. In formula, use [ ] around complex and list metal first, then ligands in alphabetical order.
    • Use abbreviations for polydentate ligands. en, ox
    • 6. Complex anions: metal ends with -ate (e.g. chromate) exceptions: Fe -> ferrate, Cu -> cuprate (Table 23.9)
  38. delta G = delta H - T(delta S)
    • Delta H  Delta S   Delta G      Result
    • -             +          -        Always spontaneous
    • -            -            -      Low T (spontaneous)
    •                           +      High T (non-spon)
    • +            +         +     Low T (non-spon)
    •                            -    High T (spontaneous)
    • +           -           +     Never spontaneous
  39. Crossover temp
    • when delta G = delta H - T(delta S) = )
    • delta H = T(detla S)
  40. Supercooling
    Lowering the temperature of a liquid below its freezing point without it becoming a solid
  41. Metastable
    System that is temporarily "trapped" in an excited energy state
  42. Many ways to compute delta Go
    • I. Use delta Go = delta Ho - T(delta So)
    • possibl to use at many temps
    • ii. Use tabulated std free energy of formations (delta Gfo)
    • Summations, mp = moles of product
    • nr = moles of reactant
    • mathematically simpler
    • can only use for rxns at 298 K
    • iii. delta Go = -RT ln K
  43. Le Chatelier's Free energy, equilibrium and Rxn Direction
    • = relative amounts of reactants or products can drive a reaction
    • ie. when not at 1 bar/1M
    • Use: delta G = delta Go + RT ln Q
    • Q = Rxn Quotient
    • Delta G is driving force of a reaction
    • At equilibrium: delta Go = -RT ln K
    • Q<K rxn procees to right: forward reaction delta G < 0
    • Q>K forward reaction has delta G >0, and reverse reaction delta G < 0.  Therefore rxn proceeds to the left
    • Q=K at equilibrium delta G=0

    Another way to determine K or delta G0
  44. Electricity
    The flow of electrons. Voltage is the pressure, current is how much flow there is.
  45. Electrochemistry
    The reactions of molecules and atoms with electrons
  46. Cell potential
    How readily two reactants will undergo electrochemical reaction
  47. Ecell
    Cell potential is measured in voltage (V)
  48. E0cell
    Standard cell potential.  All reagents 1M or 1 atm at 250C
  49. Galvanic cell, voltaic cell
    • Spontaneous chem rxns that generate electric current
    • E0cell > 0
  50. Electrolysis
    • Non-spontaneous rxns, require an electric current to produce chemical change
    • E0cell < 0
  51. Examples of electrochemistry
    Fuel cell vehicles, metal refining, metal plating, rust, etc
  52. Reduction oxidation (redox)
    Movement of electrons from one species to another
  53. Reduction (verb)
    Gain of electrons Cu2+ to Cu (s)
  54. Oxidation (verb)
    Loss of electrons Zn (s) to Zn2+
  55. Reducing agent
    Causes another species to be reduced Zn (s)
  56. Oxidizing agent
    Cases another species to be oxidized Cu2+
  57. Anode
    • Oxidation Half Rxn
    • Zn (s) -> Zn2+ (aq) + 2e-
  58. Cathode
    • Reduction Half Rxn
    • Cu2+ (aq) + 2e- -> Cu(s)
  59. Direction of electricity
    Electricity moves from Anode to Cathode
  60. Relative Strenths of Oxidzing and Reducing Agents
    • Strong oxidizing agent (positive Eo)
    • Weak oxidizing agent (negavite Eo)
    • Weak reducing agent = strong oxidizing agent
    • Strong reducing agent = weak oxidizing agent
  61. Eocell
    = Ecathode - Eanode
  62. Electrolyte
    Mixture of ions (usually aq) involved in redox rxn or carry charge
  63. Notation for a voltaic cell
    • Anode is on left
    • Cathode is on right
    • Phase boundary is denoted by |
    • Half-cell boundary (salt bridge) is denoted by ||
    • Ignore non-redox ions in electrolyte
    • If half-cell rxn has no solid species to serve as electrode, use inert electrode like Pt or graphite (Exp V)
  64. Cell Potential
    • Ecell > 0 spontaneous rxn
    • Is driving force on e-
    • units of volt: 1 V = 1 J/C (amount of energy per unit of charge)
    • -potential depends on conc of redox species,
  65. Delta G & Electrical Work
    • Amount of work done by a reaction is measured by delta G (J)
    • Delta Go = -RT ln K
  66. relationships in equations
    • delta G0 = -nFE0cell
    • delta G0 = -RT ln K
    • Eocell = (RT ln K)/(nF)
  67. Battery
    • Self contained group of voltaic cells arrganed in series
    • Spontaneous electrochemical reaction
  68. Primary Battery
    • Batteries (voltaic cell) that cannot be recharged
    • Primary lithium battery
    • Leclache (dry) cell
    • Alkaline Battery
  69. Secondary Battery
    • Batteries (voltaic cell that can be recharged
    • Lithium ion battery
    • Lead Acid Battery (car battery)
  70. Fuel Cell (flow battery)
    • Reactants (combustible fuel and oxygen) flow into battery and products leave cell
    • -Use combustion to produce electricity
    • Delta E = q + w
    • Combustion engine = q + P(delta V)
    • Fuel cell = q + welec
  71. Electrolytic cells
    Electrochemical cell that uses electrical energy to drive a nonspontaneous (Eo < 0, delta Go > 0) chemical reaction
  72. Displacement series
    Listing of reduction half rxns worst reducing agent to best (Scheme B in lab manual)
  73. Lelanche (dry) cell
    ie. Flashlight battery
    • Anode: Zn (s) -> Zn2+ + 2e-
    • Cathode: 2 MnO2 (s) + H2O (l) + 2e- _> Mn2O3 (s) + 2 OH- (aq)
    • Cell: Zn(s) + 2MnO2(s) + H2O (l) -> Zn2+(aq) + Mn2O3 (s) + 2OH- (aq)
    • Ecell = 1.55V
    • NH4Cl paste reacts consuming OH- to maintain pH but that generates NH3 (g).  Build-up of NH3 (g) minimized by formation of complex ion Zn2+ + 2NH3 (g) + 2Cl- (aq) --> [Zn(NH3)2]Cl2 (s)
    • But when current is drawn rapidly from cell, NH3 builds up near electrode, causing the voltage to drop. Also, acidic electroplyte slowly dissolves Zn (s) electrode
    • Acid dissolved Zn2+
    • Limited shelf life
    • Alkaline battery has longer shelf life...
    • Cannot recharge because products do not remain in intimate contact with electrodes
  74. Primary Lithium battery
    • High energy/mass ratio; used in watches and pacemakers
    • Anode: 3.5 Li (s) -> 3.5 Li+ + 3.5 e-
    • Cathode: AgV2O5.5(s) + 3.5 Li+ + 3.5 e- -> Li3.5AgV2O5.5
    • Cell: AgV2O5.5(s) + 3.5 Li(s) -> Li3.5AgV2O5.5(s)
    • Ecell = 3.5-4.0 V
    • Li(s) is highly reactive in water.  Therefore electrolyte must be in an organic solvent.
    • Cell can provide power for several years if used at a low rate
  75. Lead acid battery
    • Car battery,
    • Ecell =2.02 V
    • Recharge is negative, and is possible because PbSO4 sticks to electrode
    • Car battery achieves 12 V by having 6 lead acid batteries in series
    • Very low energy/mass ratio - limited use for electric cars
  76. Electrolytic Cells
    Electrochemical cell that uses electrical energy to drive a nonspontaneous (Eo < 0, delta Go >0) chemical reaction
  77. Lithium-ion battery
    • high energy/mass ratio
    • Therefore used in laptops, cell phones, etc.
    • Li(s) is highly reactive with water, the electrolyte is 1M LiPF6 in organic solvent
  78. Fuel Cell (flow battery)
    • Reactants (combustible fuel and oxygen) flow into battery and products leave cell
    • -use combustion to produce electricity
    • PEM = proton exchange membrane
    • Electrochem rxn involving gases have large Ea (overvoltage)
    • Electro-catalysts (nanoparticle Pt) lower Ea
    • delta E = q + w
    • Combustion engine = q + P(delta V) 25-40%
    • Fuel cell = q + welec 75% efficient
  79. Many transition metal compounds are
    • Highly coloured (absorb visible light)
    • Paramagnetic (unpaired electrons, interact with magnet)
  80. Ions (remember)
    • Electrons removed from valence-shell s orbital before removed from valence d orbitals
    • be able to determin electron configurations for neutrals and ions of d-block elements
  81. Coordination Compounds (complexes)
    • Consist of:
    • 1. Metal ion
    • -electropositive and a Lewis acid, an electron pair acceptor
    • 2. Ligands - molecules and/or anions with lone pairs of e-
    • -Lewis bases electron pair donors
    • -surround the metal ion
    • -bonded to the metal ion
    • 3. Counter ions
    • -ions not bonded to the central metal atom
    • -provide net zero charge for compound
  82. Coordination number
    Number of atoms bonded directly to the center metal atom (usually 2, 4, or 6)
  83. Counter-ions
    • Needed to maintain charge neutrality on the compound
    • Tells you total charge on the complex
  84. Naming Coumpounds
    • 1. Name cation first, then anion
    • 2. In naming complex, name ligands before metal. List ligands alphabetically, with anionic ligands ending in -o. Indicate number of ligands using: di-, tri, tetra-, penta-, hexa- use bis-, tris, tetrakis-... for ligands that already have the di-, tri-prefixes (ethylenediamine)
    • 3. Count up all anions and look at overall charge to determine oxidation number of metal
    • 4. Designate metal oxidation number using Roman numerals (I, II, III, IV...).
    • 5. In formula, use [] around complex and list metal first, then ligands in alphabetical order.
    • 6. Complex anions: metal ends with -ate (e.g. chromate) exceptions: Fe -> ferrate, Cu-> cuprate (table 23.9)
  85. How to get the oxidation state of a metal?
    • Look at the charge on the complex
    • Are any of the ligands negatively charged?
    • Then calculate the charge (positive) on the metal
  86. How do you tell the shape of the complex?
    • Determine its coordination number (CN)
    • Determine metal's d configuration.
    • CN = 2, Linear
    • CN = 4 Square planar(d8) or tetrahedral
    • CN = 6, Octahedral
  87. Isomers
    Compounds with same chemical formula but different properties
  88. Structural Isomers
    Atoms connected differently
  89. Coordination Isomers
    Ligand and counter-ions swapped (structural isomers)
  90. Linkage isomers
    • Different atom of a monodentate ligand bonded to metal (structural isomers)
    • ONO or NO2
  91. Geometric isomers
    Atoms arranged differently in space relative to central metal atom: cis or trans
  92. Optical Isomers (enantiomers)
    • Molecule & mirror image cannot be superimposed
    • Chemically identical properties (mirror images)
  93. Two theories that explain metal-ligand bonding:
    • Valence bond theory (VBT)
    • -assumes covalent bonding
    • -explains geometry of complexes
    • Crystal field theory (CFT)
    • -assumes ionic interactions
    • -explains color and magnetism
  94. Valence Bond theory (VBT)
    • metal - Lewis acid - accepts electrons
    • ligand - Lewis base - donated electrons
    • metal-ligand bond is fully covalent
    • equal sharing of e- pair by metal and ligand
    • bond formed by overlap f filled ligand orbital and empty metal orbital
  95. Pairing energy
    Energy required for an electron to enter a partially filled orbital
  96. geometry due to hybridized orbitals on metal
    • linear: metal orbital are sp hybridized
    • tetrahedral:metal orbitals are sp3 hybridized
    • square planar: metal orbitals are dsp2 hybridized
    • octahedral: metal orbitals are d2sp3 hybridized
    • VBT describes covalent nature of M-L bonds
    • VBT does not predict color or magnetic
  97. Crystal Field Theory (CFT)
    • CFT does not accurately describe M-L bonding
    • CFT predicts colors and magnetic properties
    • Metal-ligand bonding is considered fully ionic
    • -electrostatic attraction between metal cation (+) and negative ligands
    • -either full negative charge
    • -or negative dipole
    • But are electrons in d-orbitals of metals
    • -electrostatic repulsion between d-electrons and negative ligands
    • -results in destabilization of all d-orbitals
    • -results in splitting of the metal d orbitals
  98. Crystal Field Splitting
    • The delta crystal field splitting depends on:
    • 1. Geometry of coord cps
    • 2. Oxidation number of metal
    • 3. How focused ligand charge is
  99. Electron configuration
    • Competition between splitting energy and Epairing
    • Low spin = Epairing<splitting energy
    • High spin = Epairing > splitting energy
  100. Paramagnetic
    • Metals with unpaired electrons
    • -affected by magnets
    • -more unpaired electrons more affects by magnets
  101. Diamagnetic
    • Metals with all paired electrons
    • Unaffected by magnet
  102. Why are coordination complexes highly colored?
    • Color can be explained using Crystal Field Theory
    • Color caused by absorbance of a photon equal in energy to the difference in orbital energy
    • This occurs due to excited states, which normally returns to ground state in 10-12 seconds
    • Sometimes can undergo a reaction
    • Sometimes excess energy is emitted as light (fluorescence)-slow process (>10-9s)
  103. Color theory
    • White light = all colors
    • Color observe = colors not absorbed
    • Color characterized by wavelength
    • Ephoton = hv = h (c/wavelength)
    • On the color wheel, what you see is opposite of what wavelength they observe
  104. Tetrahedral CFT
    • 1. Different splitting pattern
    • 2. Smaller splitting energy - No orbitals directly overlap ligands, splitting energy of tetrahedral < splitting energy of octahedral
    • 3. Small splitting energy means tetrahedral complexes always weak-field (high-spin)
  105. Square Planar CFT
    • 1. Different splitting pattern
    • 2. Total splitting energy for square planar similar to splitting energy of the octahedral
  106. Hemoglobin
    oxygen carrying protein

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