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Exam 1 Practice

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Exam 1 Practice Chem 130


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Practice Problems for Exam 1
General Chemistry II
Dr. Yang
Study Guide
50 ?




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This 8 page Study Guide was uploaded by n/a on Thursday February 11, 2016. The Study Guide belongs to Chem 130 at University of Tennessee - Knoxville taught by Dr. Yang in Spring 2016. Since its upload, it has received 592 views. For similar materials see General Chemistry II in Chemistry at University of Tennessee - Knoxville.


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Date Created: 02/11/16
1.    What  is  a  manometer?  How  does  it  measure  the  pressure  of  a  sample  of  gas?         2.  Explain  how  the  ideal  gas  law  contains  within  it  the  simple  gas  laws  (show  an  example).           3.  When  a  gas  is  collected  over  water,  is  the  gas  pure?  Why  or  why  not?  How  can  the  partial   pressure  of  the  collected  gas  be  determined?           4.  The  pressure  on  top  of  Mount  Everest  (elevation  29,029  ft)  averages  about  235  mmHg.   Convert  this  pressure  to:     a)torr             c)Pa       b)psi             d)  atm       5.  A  sample  of  gas  has  an  initial  volume  of  13.9  L  at  a  pressure  of  1.22  atm.  If  the  sample  is   compressed  to  a  volume  of  10.3  L,  what  is  the  pressure?         6.  A  balloon  contains  0.158  mol  of  gas  and  has  a  volume  of  2.46  L.  If  we  add  0.113  mol  of  gas  to   the  balloon  (constant  T  and  P)  what  is  the  final  volume?         7.  A  1.0-­‐L  container  of  liquid  nitrogen  is  kept  in  a    closet  measuring  1.0  m  by  1.0  m  by  2.0  m.   Assuming  that  the  container  is  completely  full,  that  the  temperature  is  25.0°C,  and  that  the   atmospheric  pressure  is  1.0  atm,  calculate  the  percent  (by  volume)  of  air  that  is  displaced  if  all  of   the  liquid  nitrogen  evaporates.  (Liquid  nitrogen  has  a  density  of  0.807  g/mL).             8.  What  is  the  pressure  in  a  15.0  L  cylinder  filled  with  32.7  g  of  oxygen  gas  at  a  temperature  of   302  K?           9.  Use  the  molar  volume  of  a  gas  at  STP  to  determine  the  volume  (in  L)  occupied  by  33.6  g  of   neon  at  STP.           10.  A  248-­‐mL  gas  sample  has  a  mass  of  38.8  mg  at  a  pressure  of  721  mmHg  and  a  temperature   of  55°C.  What  is  the  molar  mass  of  the  gas?             11.  Use  the  molar  volume  of  a  gas  at  STP  to  calculate  the  density  (in  g/L)  of  nitrogen  gas  at  STP.           12.  A  275-­‐mL  flask  contains  pure  helium  at  a  pressure  of  752  torr.  A  second  flask  with  a  volume   of  472  mL  contains  pure  argon  at  a  pressure  of  722  torr.  If  the  two  flasks  are  connected  through   a  stopcock  and  the  stopcock  is  opened,  what  is  the  partial  pressure  of  each  gas  and  the  total   pressure?           13.  Oxygen  gas  reacts  with  powdered  aluminum  according  to  the  following  reaction:       4AL(s)  +  3O (g)  →22l O (s)   2 3                   What  volume  of 2O  gas  (in  L)  measured  at  782  mmHg  and  25°C  completely     reacts   with  53.2  g  Al?           14.  In  a  common  classroom  demonstration,  a  balloon  is  filled  with  air  and  submerged  in  liquid   nitrogen.  The  balloon  contracts  as  the  gases  within  the  balloon  cool.  Suppose  a  balloon  initially   contains  2.95  L  of  air  at  a  temperature  of  25°C  and  a  pressure  of  0.998  atm.  Calculate  the   expected  volume  of  the  balloon  upon  cooling  to  -­‐196°C  (the  boiling  point  of  liquid  nitrogen).   When  the  demonstration  is  carried  out,  the  actual  volume  of  the  balloon  decreases  to  0.61  L.   How  does  the  observed  volume  of  the  balloon  compare  to  your  calculated  value?  Explain  the   difference.             15.  Calculate  the  root  mean  square  velocity  and  kinetic  energy  of  CO,  CO ,  and  SO  at  298  K.   2 3 Which  gas  has  the  greatest  velocity?  The  greatest  kinetic  energy?  The  greatest  effusion  rate?             16.  Why  are  intermolecular  forces  generally  much  weaker  than  bonding  forces?           17.  What  is  surface  tension?  How  does  surface  tension  result  from  intermolecular  forces?  How   is  it  related  to  the  strength  of  intermolecular  forces?             18.  Arrange  these  compounds  in  order  of  increasing  boiling  point:  CH CH CH ,  CH CH Cl,   4,   3 3 3 2 CH CH OH.3   2           19.  The  human  body  obtains  915  kJ  of  energy  from  a  candy  bar.  If  this  energy  were  used  to   vaporize  water  at  100.0°C,  how  much  water  (in  L)  could  be  vaporized?  (Assume  the  density  of   water  is  1.00  g/mL).             20.  How  much  heat  (in  kJ)  is  evolved  in  converting  1.00  mol  of  steam  at  145°C  to  ice  at  -­‐50°C?   The  heat  capacity  of  steam  is  2.01  J/g*°C,  and  that  of  ice  is  2.09  J/g*°C.             21.  The  vapor  pressure  of  water  at  25°C  is  23.76  torr.  If  1.25  g  of  water  is  enclosed  in  a  1.5  L   container,  is  any  liquid  present?  If  so,  what  is  the  mass  of  the  liquid?                 22.  Liquid  nitrogen  can  be  used  as  a  cryogenic  substance  to  obtain  low  temperatures.  Under   atmospheric  pressure,  liquid  nitrogen  boils  at  77  K,  allowing  low  temperatures  to  be  reached.   However,  if  the  nitrogen  is  placed  in  a  sealed,  insulated  container  connected  to  a  vacuum   pump,  even  lower  temperatures  can  be  reached.  Why?  If  the  vacuum  pump  has  sufficient   capacity  and  is  left  on  for  an  extended  period  of  time,  the  liquid  nitrogen  to  starts  to  freeze,   Explain.                   23.  What  is  entropy?  What  role  does  entropy  play  in  the  formation  of  solutions?               24.  What  does  the  statement  like  dissolves  like  mean  with  respect  to  solution  formation?         25.  For  each  compound,  would  you  expect  greater  solubility  in  water  or  in  hexane?  Indicate  the   kinds  of  intermolecular  forces  that  occur  between  the  solute  and  the  solvent  in  which  the   molecule  is  most  soluble.       a)  glucose           c)  dimethyl  ether       b)  naphthalene         d)  alanine       26.  Calculate  the  mass  of  nitrogen  dissolved  at  room  temperature  in  an  80.0  L  home  aquarium.   Assume  a  total  pressure  of  1.0  atm  and  a  mole  fraction  for  nitrogen  of  0.78.                 27.  To  what  volume  should  you  dilute  50.0  mL  of  a  5.00  M  KI  solution  so  that  25.0  mL  of  the   diluted  solution  contains  3.05  g  of  KI?                   28.  A  solution  contains  a  mixture  of  pentane  and  hexane  at  room  temperature.  The  solution  has   a  vapor  pressure  of  258  torr.  Pure  pentane  and  hexane  have  vapor  pressures  of  425  torr  and   151  torr,  respectively  at  room  temperature.  What  is  the  mole  fraction  composition  of  the   mixture?  (Assume  ideal  behavior.)                     29.  Calculate  the  freezing  point  and  melting  point  of  a  solution  containing  7.55  g  of  ethylene   glycol  (C H O )  in  85.7  mL  of  ethanol.  Ethanol  has  a  density  of  0.789  g/cm .   3 2 6 2                       30.  The  density  of  a  0.438  M  solution  of  potassium  chromate  at  298  K  is  1.063  g/mL.  Calculate   the  vapor  pressure  of  water  above  the  solution.  The  vapor  pressure  of  pure  water  at  this   temperature  is  0.0313  atm.  (Assume  complete  dissociation  of  the  solute.)              


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