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Mason - CHEM 211 - CHEM 211, Week 7 Notes - Class Notes

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Mason - CHEM 211 - CHEM 211, Week 7 Notes - Class Notes

This preview shows pages 1 - 2 of a 3 page document. to view the rest of the content CHEM 211  MW 10:30-11:45 AM  Prof. Gregory Foster  Page 1 of 3     Notes for Week of 10/16  I.  Dalton’s law of partial pressure  a.  Ideal gases behave the same.
b.  P(total) = P1+P2+…+Pn
c.  P1= n1RT/V
i.  Ptotal = (n1+…+nn)RT/V  d.  P1=n1RT/V / ntotal   i.  P1/ntotal = (n1/nt)(RT/V)  ii.  (n1/nt)(ntRT/V)  iii.  P1= n1/nt *Ptotal  e.  A capital X is concentration (mol fraction) = n1/nt  i.  P1=X1Ptotal  f.  Sample Question  i.  A 335.0 mL container holds 0.146 g o Ne and an unknown amount of Ar  at 35 degrees C. The total P is 626 mmHg. Calculate the number of mol of
Ar present.
1.  Ntotal = n(neon)+n(argon)  a.  .146 g neon * 1 mol/ 20.18 g Ne = .00723 mol neon  2.  Ntotal = PtotalV/RT  a.  (0.824atm)(0.335 L)/( .0821 Latm/molK)(308K) = 0.0116  mol total  3.  0.0116 mol total = 0.00723 mol Neon + n(argon)
4.  N(argon) =  0.00437 mol
II.  Ideal gas law and stoichiometry  a.  How many grams of P4 react with 35.5 L of O2 at STP to form P4O10  i.  P4 +5 O2  P4O10  ii.  We need to find the mol of everything  iii.  nO2 = PV/RT = (1atm)(35.5 L)/ (.0821 Latm/molK)(273K) = 1.58 mol O2  iv.  1.58 mol O2 * 1 mol P4/ 5 mol O2 * 123.88 gP4/1 mol P4 = 39.2 g P4  III.  Kinetic Molecular Theory (KMT)  a.  Particle volume in a system is very low – gases represent a very small fraction of  the volume of a system  b.  Particle motion – gas particles move in a straight line; until there are collisions  between particles, then they move off in a straight line until they have another
collision with either a particle or a container wall
c.  Particle collisions – are elastic; kinetic energy (1/2mv^2) is conserved in a  collision  d.  As you raise temperature gases move faster and the total number of velocities  increases as well    i.  On a distribution we call this degrees of freedom  e.  Ek-bar = 1/2mv^2
f.  Particle velocity – Ek-bar =cTemperature
i.  C= 3*R/2*6.02e23) CHEM 211  MW 10:30-11:45 AM  Prof. Gregory Foster  Page 2 of 3     g.   V = sqrt((3RT)/(Mm))
h.  What is rms veleocity of O2 at STP?
i.  Vrms =sqrt((3*8.314J/molK *273K)(0.032 kg/mol)) =461 m/s  IV.  Effusion and diffusion  a.  Effusion – is how we purify gas by passing gases through filters  V.  Deviations from the ideal gas law  a.  Ideal gas law works when total pressure is fairly low
b.  Van  der Waal equation helps calculate actual gas laws when IGL doesn’t work
c.  As pressure increases the effect of particle volume predominates
d.  As pressure decreases than the effect of interparticle attraction predominates
VI.  Thermochemistry  a.  Kinetic and potential energy  i.  deltaE = internal energy = Efinal – Einitial  ii.  deltaE = q + w  1.  q = heat in (J)(kgm^2/s^2)
2.  w = work is typically expressed as  PdeltaV
3.  heat content is easy to measure with a thermometer
iii.  a change in energy is a change in the total internal energy  iv.  a chmical system is everything that isn’t the surroundings. Typically  isolated  1.  a negative q shows that heat is leaving the system and going to the  surroundings  a.  a positive q shows that heat is entering the system  2.  if we have a positive delta V then the system is doing work on the  surroundings  a.  if volume expands or gets bigger and doing work on the  surrounding then it is losing energy to the surroundings and
is therefore written as  -w
v.  sample problem   1.  a reaction profuces a gas and the volume of container increases  from 125 to 652 mL against a pressure of 988 torr. Calculate the
work done (J)
a.  w = -PdeltaV
b.  988 torr * 1 atm/760 mmHg 8 101,325 Pa/ 1 atm
c.  1.32e5 Pa
d.  deltaV =Vf-Vi = 0.625L – 0.125L = 0.500L
e.  0.500 L* 1m^3/1000L = 5.0e-4m^3
f.  W= -(1.32e5 Pa)(5.0e-4m^3) = 66 Pa-m^3 = -66 J
b.  Heat and work  i.  Higher heat with more kinetic energy  ii.  Heat is the transfer of thermal kinetic energy from one particle to another  iii.  Works by vibration = higher motion = more energy = more heat  c.  First law of thermodynamics

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