General Chemistry

  1. charge of 1 e-
    1.6 x 10-19 C
  2. isotopes (# of protons and neutrons)
    • same # of protons
    • different # of neutrons
  3. metals = oxidizing or reducing agents?
    reducing agents

    tend to lose e-
  4. elements of same family or period are most similar
  5. quantum mechanics: what does n and l represent?
    • n = energy level
    • l = shape of orbital
  6. naming (order)
    • - ide
    • - ite
    • -ate

    • "hypo" = first
    • "per" = last
  7. quantum mechanics: increase in n does what?
    increases size and energy of e- orbital
  8. quantum mechanics: e- in the same orbital share what common characters?
    • same 1st 3 numbers, different 4th:
    • same shell (n), subshell (l), orbital
    • different spin (+1/2) or (-1/2)
  9. what is electron affinity?
    amt. of energy released when an e- is gained
  10. what is electronegativity?
    ability of atom to draw in e- from a bond
  11. periodic trend: electron affinity
    increases right to left, bottom to top

    diagonally left and right
  12. periodic trend: electronegativity
    increases left to right, bottom to top
  13. periodic trend: acidity
    increases left to right, top to bottom

    (right diagonal down)
  14. noble gases have exclusively what intermolecular bonds?
    van der Waals
  15. Are larger or smaller atoms more polarizable?
    larger atoms - further away
  16. van der waals interactions are:
    temporary dipole-dipole moment
  17. relative strengths of intermolecular forces
    • H bonding
    • dipole-dipole
    • van der Waals
  18. STP (temperature and pressure)
    • 0oC (273K)
    • 1 atm (760 torr)
  19. At STP, how much volume does 1 mole of a gas occupy?
  20. describe ideal gases
    • 1. zero volume
    • 2. no forces other than repulsive in collision
    • 3. elastic collisions
    • 4. avg. KE proportional to temperature
  21. are intermolecular forces taken into account in ideal gases?
  22. PV = nRT (R value)
    8.314 J
  23. ideal gas: KE and temperature
    avg. KE is proportional to temperature
  24. what is greater, ideal pressure/volume or real pressure/volume
  25. what deviates from "ideal"
    when molecules are close together
  26. ideal temperature and pressure
    high temperature and low pressure

    (so molecules are far apart)
  27. what are more ideal: monoatomic gases or diatomic gases (H2, Br2 etc.)
  28. noble gases at STP are what kind of gases?
  29. effusion
    pinhole - spreading from high to low pressure
  30. which gases will find pinhole easier?
    gases with higher rms velocity
  31. diffusion
    spreading into another gas or empty space
  32. Partial pressure equation
    PA = XA times Ptotal
  33. avg. KE of gases in a gas mixture
    KE = 3/2 RT
  34. decrease KE of a gas will do what to PE?
    increase PE

    molecules bind together more
  35. will KE be the same throughout a gas mixture?
    no, only their average kinetic energies
  36. why will KE vary from molecule to molecule?
    they have different molecular weights
  37. equation of KE of gases in a mixture
    velocity1/velocity2 = sq. root (mass2/mass1)

    reversed subscripts
  38. effusion and diffusion rates
    reversed subscripts

    inversely proportional to sq. root of their mass
  39. which gas will diffuse/effuse faster O2 or H2?
    H2 because it's lighter
  40. as gas with strong or weak IM forces will exert more pressure on a container?
    weak intermolecular forces
  41. Kinetics: rate of a reaction increases when:

    [reactants] and temperature
    [reactants] increases

    temperature increases
  42. half life of 1st order reaction
  43. Kinetics v. Thermodynamics
    • kinetics = rate (getting to state)
    • thermodynamics = G, S, H (prop. of this state)
  44. heat transfer (3 forms)
    • conduction
    • convection
    • radiation
  45. how does conduction transfer heat?
    molecular collisions
  46. slab
    • conducts heat from hot to cold reservoir
    • (higher energy to lower energy)
  47. conducting capability of slab and length
    • doesn't matter -
    • same regardless of length
  48. what about temperature difference between slabs?
    increase distance = increase temperature difference
  49. Is rate of flow across slabs constant or varies?
  50. convection transfers heat how?
    fluid movement
  51. radiation transfers heat how?
    electromagnetic waves
  52. What is work?
    energy transfer that is not heat
  53. Work (pressure and volume)
    W = -P x change in V
  54. work at constant volume

    no work is done
  55. Total energy of a system:
    E = q + W
  56. total energy at constant volume
    since no work, E = q
  57. work at constant temperature
    W = -q
  58. total energy at constant temperature
    E = 0 because E is proportional to T
  59. energy flow into and out of the system
    • into (-)
    • out of (+)
  60. How is enthalpy calculated? What is the formula?
    H = change U + P changeV

    H = E
  61. What is enthalpy of formation?
    change in enthalpy when a compound is formed from raw elements
  62. enthalpy at constant P
    H = change U + P change in V

    H = W + U
  63. element in standard state at 25oC
    H = 0 J/mol
  64. endothermic v. exothermic and delta H
    • endothermic (+) - absorbs heat
    • exothermic (-) - releases heat
  65. energy of activation greater for endothermic or exothermic?
    greater in endothermic
  66. what elements are Hfo = 0?
    standard state elements - O2 H2
  67. entropy of universe must always be positive or negative?
  68. equation for entropy
    delta S = Q/T

    energy transfer/temperature
  69. entropy change between hot and cold reservoir: positive or negative?
    hot = negative (energy is leaving system)

    cold = positive (energy is entering system)
  70. entropy for any isolated system that is irreversible must be....

    entropy of universe must always be > 0
  71. increase # of moles of gas does what to entropy?
    increases entropy
  72. increase in temperature = what to entropy
    increase in entropy
  73. reaction at equilibrium will maximize....
    universal entropy (not entropy of system)
  74. spontaneous = positive or negative entropy?
    positive (more disorder)
  75. exergonic and endergonic
    • exergonic = spontaneous
    • endergonic = non-spontaneous
  76. delta G and spontaneity
    • (-) = spontaneous
    • (+) = non-spontaneous
  77. What is Gibbs Free Energy?
    energy available to do work
  78. equation for Gibbs Free Energy
    delta G = delta H - Tdelta S
  79. more disordered (greater S) does what to Free energy
    negative free energy = spontaneous

    because universe tends towards greater disorderedness
  80. G and K
    delta G = -RT lnK
  81. G = 0 then K
    K = 1
  82. what does K = 1 indicate?
    products and reactants are equal
  83. G < 0 and G > 0
    • (-) G = K > 1
    • (+) G = K < 1
  84. K < 1
    reactants are favored
  85. K > 1
    products are favored
  86. Molarity
  87. molality
  88. 1L of water = how many kilograms?
    1L = 1kg of water
  89. 1L water = how many moles?
    55.5 moles
  90. solution mixture v. separated pure substances (G and S)
    solution = more disordered = inc. entropy = decrease G
  91. Ksp
    • [products]coefficients /
    • [reactants]coefficients
  92. pure solids and liquids
    leave them out
  93. what is solubility? (measured)
    # moles/ L
  94. how does pressure affect solubility of gases?
    increases solubility
  95. how does temperature affect solubility of salts
    increases solubility of salts
  96. how does temperature affects solubility of gases?
    decreases solubility
  97. What is heat capacity?
    measure of energy change needed to change temperature of substance
  98. constant volume:
    no work (W=P change in V)
  99. work at constant pressure... why?
    yes because substance can expand, so substance absorbs energy
  100. heat capacity formula
    q = C x change in T
  101. heat capacity per unit mass
    q = mc change T
  102. are heat capacities different according to phase?
  103. specific heat of water
    1 cal/goC
  104. joules/calorie conversion
    1 calorie = 4 joules
  105. when a substance releases heat, temperature or pressure change?
    yes, but either temperature or heat, not both
  106. increasing non-volatile solute does what to vapor pressure?
    decreases vapor pressure
  107. increasing non-volatile solute does what to boiling point
    decreases boiling point
  108. osmotic pressure formula

    Molarity x 0.8 x T x i
  109. Free Energy Equation
    delta G = H - TS
  110. entropy at low temperatures
    not much influence
  111. ions dissolved in solution are called

    able to conduct electricity
  112. is water a conductor of electricity?
    yes, polar conductor

    unless it contains electrolytes
  113. 0.2 moles of NaCl will dissociate how?

    will produce 0.2 moles of Na+ 0.2 mol of Cl-
  114. vapor pressure: what holds down the molecules?
    intermolecular forces
  115. how can molecules break intermolecular forces?
    have enough KE (for a gas)
  116. if molecules are rising above to space, how can they be in equilibrium
    force some mol. to crash back down

    eq. = same # leaving and returning
  117. delta Hsolution (+) and (-)
    negative = stronger bonds are formed

    positive = weaker bonds are formed
  118. stronger bonds and weaker bonds affecting vapor pressure
    strong = lowers vapor pressure

    weaker = raises vapor pressure
  119. saturated solution
    concentration of dissolved salt reaches maximum
  120. vapor pressure and boiling point
    increased vapor pressure = decrease in boiling point
  121. Heat of transition (phase change)
    magnitude is related to.....
    strength of intermolecular forces
  122. delta H fusion
  123. delta H vaporization
  124. are melting and boiling endothermic or exothermic

    heat is being added, temperature increases
  125. sublimation
    solid to gas
  126. deposition
    gas to solid
  127. fusion
    solid to liquid (melting
  128. crystalization
    liquid to solid (freezing)
  129. triple point
    exists in equilibrium as a solid, liquid, gas
  130. critical temperature
    temp. where a substance cannot be liquefied
  131. critical pressure
    pressure required for liquidification at critical temp.
  132. critical point
    critical temperature and critical temperature
  133. 1 atm in phase diagram
    boiling point
  134. definition of Lewis acid/base:
    accepts/donates e-
  135. definition of Bronsted/Lowry acid and base
    donate/accept H+
  136. amphoteric
    acts as either acid or base
  137. example of amphoteric substance
    diprotic acid
  138. polyprotic
    acids that can donate more than one proton
  139. Ca(OH)2
    all strong bases
  140. strong acids and bases in water....
    completely dissociate
  141. F - Cl - Br - I

    decreasing polarity
  142. F - Cl - Br - I

    bond strenght
    decreasing bond strengh
  143. F - Cl - Br - I

    increasing acidity
  144. conjugate base: more oxygens...
    stronger acids
  145. why do more oxygens make stronger acids?
    O draws e- to one side, increasing polarity
  146. periodic table: acidic and basic
    • left = basic
    • right = acidic
  147. periodic trend for acidity
    increases right and down
  148. hydrides are what kind of compounds?
    binary compounds = only 2 elements
  149. pH equation
    pH = -log[H+]
  150. pOH equation
    pOH = -log[OH-]
  151. pH + pOH add to...
    pH + pOH = 14
  152. [H+] and [OH-] of strong acids/bases in solution
    strong acids and bases dissociate completely in solution -

    [H+] and [OH-] are the same as original conentrations
  153. weak acids and bases concentrations [H+] and [OH-]
    requires Ka
  154. What is the pH of 1M of HCl?
    pH = 0

    because [H+] = 1 and 100 = 1
  155. What is value of Ka?
    Ka = [H+][A-] / [HA]
  156. high Ka value means what
    strong acid
  157. pKa equation
    pKa = -log[Ka]
  158. small pKa means what value of Ka
    small pKa = large Ka

    small pKa = strong acid
  159. neutralization
    acids and bases combine exothermically to from water + salt
  160. dilution formula
    MinitialVinitial = MfinalVfinal
  161. how do you determine pH of salt solution?
    look at conjugates of ion
  162. ex: Na Cl dissociation neutralization
    conjugates = NaOH and HCl

    both strong, so Na+ and Cl- is neutral
  163. buffering region of titration curve
    flat plateu
  164. equivalence point
    all acid has been converted to base

    100% A-
  165. equivalence point of titration of strong acid by strong base
  166. equivalence point: stronger acid than base
    EP < 7
  167. equivalence point: stronger base than acid
    pH > 7
  168. half equivalence point
    [acid] = [base]
  169. polyprotic titrations
    have more than 1 EP and more than 1 1/2 EP
  170. polyprotic titrations: when does second proton begin to dissociate?
    once the 1st proton has completely dissociated
  171. buffered solution
    no change in pH upon addition of acid or base
  172. At what point in the titration curve is the solution buffered?
    1/2 equivalence point

    because [HA] = [A-]
  173. Henderson-Hasselbalck
    pH = pKa + log[A-]/[HA]
  174. pH at half-equivalence point
    [HA] = [A-] so pH = pKa
  175. how do you make buffer solutions?
    mix equal amounts of acid with conjugate base
  176. indicators (colors)
    basic = red to blue

    acidic = blue to red
  177. endpoint in titration
    when indicator changes color
  178. oxidation
    loss of e-
  179. reduction
    gain of e-
  180. the lion says.....
    LEO the lion says GER
  181. oxidation state of F, H, and O
    • F = -1
    • H = +1
    • O = -2
  182. is water and oxidizing or reducing agent?
    poor oxidizing and reducing agent
  183. balancing redox rxns:
    • balance other elements
    • H2O - balance O
    • H+ - balance H
    • e- - balance charge
  184. energy transfer in galvanic cell
    chemical energy to electrical energy
  185. T - E - I - E' - T
    • T = terminals
    • E = electrodes
    • I = ionic conductor
  186. what is the salt brigade?
    ionic conducting phase
  187. what does ionic conductor (salt bridge) do?
    carries current (in the form of ions)
  188. emf
    electric potential difference from T to T
  189. anode and cathode (charge)
    • anode = (-)
    • cathode = (+)
  190. mnemic for redox reactions

    reduction cathode - anode oxidation
  191. when is it spontaneous?
    positive cell potential (E)
  192. what is cell potential? (emf)
    potential difference between terminals when they're connected
  193. what does connecting terminals do to potential difference
    reduces potential difference
  194. cell diagram
    Anode l Anoic solution ll Cathodic solution l Cathode
  195. free energy and chemical energy equation
    delta F = -nFE
  196. what does n represent [G = -nFE]
    number of moles
  197. F
    Faraday's constant:

    100,000 C/mol
  198. Nernst Equation
    E = Eo (RT/nF) lnQ
  199. What is the cell potential at equilibrium?
    E = 0

    because no rxn is favored
  200. concentration cell
    identical electrodes whose half cells have different ion concentation
  201. current flows toward
    toward greater entropy
  202. electrolytic cells
    require force by outside power source
  203. Cell potential for H2
    2H+ + 2e- > H2 E = 0
  204. Heisenburg uncertainty principle:
    two unknowns
    position and momentum
  205. when e- fall from higher energy state to lower energy state
    photon is released

    (energy is released in form of a photon)
  206. what is the energy of this photon?
    E = hf
  207. what happens when a photon is absorbed?
    e- is bumped to a higher energy state
Card Set
General Chemistry
General Chemistry MCAT