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CHEM 121 Chemistry Assignment

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CHEM 121 Chemistry Assignment

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Course Code: CHEM 121
University: University Of British Columbia

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Country: Canada

1. If the temperature of a 1.00 L sample of gas is changed from 30.0?C to 20.0?C while the pressure is held constant, what is the final volume of the sample in liters?
2. A container of pure helium has a volume of 200.0 mL, a temperature of 50.0?C, and a pressure of 2.00 atm. How many moles of helium are present?

All gases becomes fluids at certain temperatures and pressure
As the absolute temperature increases, the average kinetic energy of the molecules increases. An increase in temperature increases the rate of collision of the particles.
Low temperatures and high pressures. The deviations of actual gases often agree with the predictions of the ideal gas equation to a limit of 5% at normal pressures and temperatures.
The entire above are correct
For the choice C, the kinetic energy refers to the speed of molecules in a given sample which is often fully dependent on the temperatures. The real molecular composition of the chemical is not of interest with kinetic energy. Should both chemical be at the same temperature then their molecules are moving at the same speed and thus their kinetic energy being identical.
P1V1=P2V2; V2=P1V1/P2
=1×293/303=0.967 L
=0.015 moles
=0.237 moles
Molar Mass=10/0.237=42.19 g/mol
N2 (g) + 3H2 (g) ® 2NH3 (g)
Moles ratio of N: H=1:3
Volume of N=1×20/3=6.667 L
=426.56 mL=4.3×102 mL
2Al(s) + 3H2SO4 ® Al2 (SO4)3 + 3H2
Mole ratio of Al: H2=2:3
Moles of H2 produced=3×0.5/2=0.75 moles
Volume of H2 produced
1 mole=22400 mL
0.75 moles=0.75×22400/1=16800 mL
=2.75 atm
=0.00143 moles
Molar Mass=0.1/0.00143=69.93g
Molecular Formula=n×Empirical formula
Molecular Formula= (CH2)5
=28.485 mL
V=nRT/P; n=m/MM
m/V= (32×0.85)/ (0.0821×373)
Mole ratio of CaC2:C2H2=1:1
Moles of C2H2
=10*20/ (0.0821*293)
=8.3147 moles
Moles of CaC2=8.3147 moles
Mass of CaC2=8.3147×64
Moles of N=63.64/14=4.5457 moles
Moles of O=36.36/16.01=2.2711 moles
Simplest Form of the moles
Empirical formula=N2O
Molecular Mass=
196 mL=0.385 g
22400 mL=22400×0.385/196=44 g
Molecular Formula=n× Empirical formula
Molecular Formula= (N2O) 1= N2O
Rate of effusion NOF/HBR= (M.W.HBR/M.W. NOF) 1/2= (80.912/49.005)1/2
The volume of the gas particles is much smaller than the distance between the gas particles. At relatively high pressure, the values of PV/RT tend to be lower than the ideal and at extremely high pressure PV/RT values are often greater than ideal.
Relationship between Hg and atm
1 atm=30 Hg
19 Hg=19/30=0.633 atm
T2=P2V2T1/P1V1= (2.50×435×355)/ (1.88×28.5)
Final temperature=720.53K
Moles of the halogen, n=PV/RT
=1.41*0.109/ (0.0821×398) =0.004706 moles
Molar Mass=0.334/0.004706=71.0 g/mol
This corresponds to the molar mass of Cl2
PV= (m/MM)/RT
MM=mRT/PV=0.465×0.0821×298/ (1.22×0.245)
=38.0 g/mol
No. of atoms
100 g contains 30.45g N and 69.55 g O
Moles of N=30.45/14.01=2.173 moles
Moles of O=69.55/16.014.347 moles
Simplest form of the Moles
Empirical formula=NO2
n=PV/RT= (1.02×0.389)/ (0.0821×273)
=0.0177 moles
Molar Mass=1.63/0.0177=92.1 g/mol
Molecular Formula =n× Empirical formula
Molecular Formula= (NO2)2= N2O4
50.0 g KClO3=50.0 gKClO3/122.55g KClO3
=0.40799 moles
2 moles of KClO3 will generate 3 moles of O2
Hence, 0.40799 moles of KClO3
0.612×0.40799×2/3 moles of O2
At STP, 1 mole of gas=22.4 L
Set up ratio
0.612 moles of O2/x litres=1 mole/22.4L
x=13.7 L of O2
All molecules have kinetic energy. The heaviest will hence move the most slowly  
Alezi, D., Belmabkhout, Y., Suyetin, M., Bhatt, P. M., Weselin?ski, ?. J., Solovyeva, V., … & Eddaoudi, M. (2015). MOF crystal chemistry paving the way to gas storage needs: aluminum-based soc-MOF for CH4, O2, and CO2 storage. Journal of the American Chemical Society, 137(41), 13308-13318
Atkins, P., De Paula, J., & Keeler, J. (2018). Atkins’ physical chemistry. Oxford university press
Chang, R., & Overby, J. (2000). General chemistry: the essential concepts. Mc Graw Hill
Haynes, W. M. (2014). CRC handbook of chemistry and physics. CRC press
Silberberg, M. S. (2007). Principles of general chemistry (p. 29). New York: McGraw-Hill Higher Education

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