The following readings were taken during the determination of the molar mass of a gas by direct weighing.?
92170
The ideal gas law is another tool that we can use to determine the molar mass of a gaseous substance. To determine a molar mass, we need the mass of a given sample and also the number of moles in that sample. You recognize immediately that given the volume, temperature, and pressure of a gaseous sabstance, you can calculate the number of moles.\n\nfirst:\n \n V = Volume = 987.56g / 1g/cm^3 = 987.56cm^3 (1 liter/ 1000cm^3) = .987 liter\n\n T = temperature = 23 + 273 =  296K\n\n P = pressure = 97.7 kPa (  1 atm / 101.3 kPa ) = .964 atm\n \n r = .0821 L'atm/mol.K\n \n n = mole = ?\n     \nThe ideal gas law equation:\n \n   PV  =  nRT\n\nsolve for  n\n\n   n = PV /RT = .964*.98756 / .0821*296 = .039 mol \n\n The molar is simply the mass divided by the number of moles:\n\n  mass of gas in flask =  187.942g - 183.257g = 4.685g\n\nso, the answer is:\n\n            4.685g / .039 mol = 120.1 g/mol
43865
The molecular mass (abbreviated MM) of a substance (less accurately called molecular weight and abbreviated as MW) is the mass of one molecule of that substance, relative to the unified atomic mass unit u (equal to 1/12 the mass of one atom of carbon-12). Due to this relativity, the molecular mass of a substance is commonly referred to as the relative molecular mass, and abbreviated to Mr.\n\nThe molecular mass can be calculated as the sum of the atomic masses of all the atoms of any one molecule. The molecular mass can also be measured directly using mass spectrometry. In mass spectrometry, the molecular mass of a small molecule (< ~200 atoms of a given element) is exact, that is, the sum of the most common isotopes of each element. Of larger molecules, it is average, or calculated using the elemental mass number or the periodic table, since there is likely to be a statistical distribution of atoms representing the isotopes throughout the molecule.\n\nThe molar mass of a substance is numerically equal to the molecular mass, but expressed in mass units per mole, usually as g/mol (grams per mole).\n\nExample: \nThe atomic mass of hydrogen is 1.00784 u and that of oxygen is 15.9994 u; therefore, the molecular mass of water with formula H2O is (2 × 1.00784 u) + 15.9994 u = 18.01508 u. Therefore, one mole of water weighs 18.01508 grams. However, the exact mass of hydrogen-1 (the most common isotope) is 1.00783, and the exact mass of oxygen-16 (the most common isotope) is 15.9949, so the mass of the most common single molecule of water is 18.0105 u. \nMolecular mass or molar mass are used in stoichiometry calculations.\n\nSince molecules are created by chemical reactions, not nuclear reactions, a molecule's molecular mass exactly equals the sum of the atomic masses of its constituent atoms.\n\nThe gram-molecular weight is the molecular weight of a chemical compound expressed in grams. Thus the molecular weight of calcium carbonate is 100 therefore the gram-molecular weight is 100 g. The Gram Molecular Weight is frequently used, particularly for calculating the concentration of solutions. Thus a molar solution will contain the gram molecular weight of a chemical dissolved to produce one litre of solution.\n\nHandy Little Calculator:\nhttp://www.chemie.fu-berlin.de/cgi-bin/molform
700533
The -891 is the molar value (for 1 mole) of methane.  Change 2 g to moles and use a ration.  In the second use PV=nRT to find the moles and use a ratio.
87793
44.00995 grams per mole
22880
There's a formula for volumes of gasses:\n\nPV = nRT.\n\nIf I remember right, P = pressure, V = volume, n = number of moles, T = temperature, and R = molar density, or something molar that is specific to the gas.
952068
a) 0.01 mol\n\n1 mole of gas, at standard tremperature and pressure occupies 22.4 l.  You have exactly 1/100 of that, therefore 1/100 of a mole.
770958
1/1000 x molar mass (grams/litre) and solute in 1 litre
365485
I assume you mean molecular weight.\n\nSince carbon weights 12 grams, and oxygen weighs 16:\n12+(16*2) = 44 grams.\n\nYou have to give the amount of CO^2, in ML to get Moles.\n\nSo generically speaking, 1 mol of CO^2 weighs 44 grams, 2 mol weights 88 grams, and so on.\n\nTo calculate the amount of atoms in it, you would have to times it with Avogadro's number. (6.0221415 × 10)
322112
That's the definition of the Molar mass, yes.
1153522
The quantities you have listed are in fact related via the ideal gas law.  So, assuming you aren't in huge temperature or pressure extremes, you can use the ideal gas law with decent results.\n\nIdeal Gas Law\nPV = nRT\n\nwhere P = pressure (N/m^2)\nV = volume (m^3)\nn = number of moles of gas\nR = universal gas constant = 8.3145 J/(mol*K)\nT = temp (in Kelvins)\n\nSo, by looking at the equation you can not convert air pressure directly to volume without also knowing the temperature and number of moles of gas.
252970
Argentometry refers to quantitative analysis using silver (remember the symbol for silver is Ag comes from the Latin word "argentum" meaning "silver") For example, silver nitrate solution could be used to determine the concentration of hydrochloric acid by precipitating the chloride as silver chloride, then drying and weighing it. Silver's high atomic mass is useful here to give high precision.
895405
use dalton's law to get the pressure of CH4(total pressure - H2 pressure)\nthen use ideal gas law to get the moles of CH4(PV=nRT)\n*watch your units*\nchange moles of CH4 to grams of CH4(moles x molar mass)
727216
use beers law to figure out the concentration and then take that to mole ratio
227577
Molar mass is 61.97894 g/mol.
839352
If tha atomic mass of M is A, then the molecular weight is MW=2*A+60\n\nThe concentration of the M2CO3 solution is C=mole/V and since mole=mass/MW, C=mass/(MW*V)\n\nThen you do a titration for which we assume complete reaction and not reaching only the first end-point, so the reaction is\n\nM2CO3 + 2HCl -> 2MCl + CO2 + H2O\n\nTherefore 2*mole M2CO3=mole HCl =>\n2*C1*V1= C2*V2=>\n2*V1*mass/(MW*V) = C2*V2 =>\nMW= 2*V1*mass/(C2*V2*V)=\n=2*0.025*7.46/(0.1*0.027*1)\n=138\n\nand MW=2*A+60=138=> 2A=78 => A=39\n\nSo M, is potassium (K)
205960
1.  convert torr to atm\n800/760 = 1.05 atm\nconvert Celsius to Kelvin\n25+273 = 298 K\nplug values into PV=nRT\n(1.05)(V)=(1 mol)(0.0821)(298)\nV = 23.3 L\nfinally... calculate molar mass ammonia = 17.04 g/mol\nand plug values into density equation D = m/V\n17.04/23.3\nD = 0.73 g/L\n\n2.  Similar to above, but just rearrange the density equation.   Use PV=nRT to find volume again (don't forget to convert units).  Then plug this volume into your density equation.  You'll need to cross-multiply to solve for "m".\n\n3.  Plug the volume, pressure, and temperature into PV=nRT and solve for "n" which is the number of moles (don't forget to convert your units).  Next, find the molar mass of the gas by dividing the 10.3 grams by "n".  You should be able to compare this to your periodic table to determine which element it matches.
28731
The molecular mass of a substance (less accurately called molecular weight and abbreviated as MW) is the mass of one molecule of that substance, relative to the unified atomic mass unit u (equal to 1/12 the mass of one atom of carbon-12).\n\nThe molecular mass can be calculated as the sum of the atomic masses of all the atoms of any one molecule. The molecular mass can also be measured directly using mass spectrometry. In mass spectrometry, the molecular mass of a small molecule (< ~200 atoms of a given element) is exact, that is, the sum of the most common isotopes of each element. Of larger molecules, it is average, or calculated using the elemental mass number or the periodic table, since there is likely to be a statistical distribution of atoms representing the isotopes throughout the molecule.