Molecular Mass

# How do you find the molecular mass of a compound? ## Related Questions

the relative molecular mass or molecular weight of a compound is the mass of a molecule of the compound relative to the mass of a carbon atom taken as exactly 12.

This is impossible. In order to find this you must know the elements in the compound

To find the molecular weight (also called the molecular mass or molar mass) of a compound or chemical, you need two things: the molecular formula of the compound and a periodic table. The molecular formula tells you how many atoms of each element is present in the compound. To find the molecular weight, just add up the atomic weights of each element present in the compound, being sure to multiply by the number of times that atom appears. The atomic weight of each element is found on a periodic table.

If you know the molar mass of the compound, you have to calculate the mass of the empirical formula and divide the molar mass of the compound by the mass of the empirical formula in order to find the ratio between the molecular formula and the empirical formula. Then multiply all the atoms by this ratio to find the molecular formula!

Molecular mass and formula mass represent the mass of an individual molecule or formula unit of a compound.

The molecular formula mass of this compound is 60.0 amu. The subscripts in the actual molecular formula are 2,4,2.

The molecular formula mass of this compound is 240 amu. The subscripts in the actual molecular formula 8,16,8.

The term formula mass is generally defined as the mass of a unit cell in an ionic compound. Molecular compounds are just defined in terms of molecular mass.

You start with the mass of the compound. This you divide by the total molecular mass of the compound (the atomic mass of 1 molecule) and you multiply this by the amount of atoms 1 molecule consists of.

Find the molar mass of CH2O, which is roughly 30amus Then divide the molecular mass of 180.15amus by 30 = 6 Multiply each element by 6 to give C6H12O6

The molecular mass of a compound is the sum of the chemical elements weights contained in the molecule.

The molecular mass of ethyl group is 29. The atomic mass of chlorine is 35.5. So, the molar mass of this compound is 64.5u.

The molecular formula is the same as the empirical formula, NO2. The compound NO2 has a molar mass of 46g/mol, so the empirical and molecular formulas are the same.

86.18 g/mol, from Wikipedia http://en.wikipedia.org/wiki/Hexane Btw, The molecular molecular weight of a compound is the same thing as molar mass. Wikipedia has molecular weight for almost every compound in the bar/table on the right side of the webpage for the compound, next the label "molar mass". 86.18 g/mol, from Wikipedia http://en.wikipedia.org/wiki/Hexane Btw, The molecular molecular weight of a compound is the same thing as molar mass. Wikipedia has molecular weight for almost every compound in the bar/table on the right side of the webpage for the compound, next the label "molar mass".

The atomic mass of I is 127. The atomic mass of F is 19. Therefore the molecular mass of the compound is 127+5x19=214u.

This compound is unstable. molar mass of it is 173gmol-1.

Hydrochloric Acid is a Compound. Molecular formula: HCl Molecular Mass: 36.46 g/mol

The molecular mass of a compound is the sum oh the atomic weights of the elements contained in the molecule of this compound.

For a compound AB:Concentration of A (%) = Atomic weight of A x 100/Molecular mass of AB

The formula of NO2 has a molecular weight of 46 g/mol. Your compound has a molecular weight of 92 g/mol. As you can see the molecular weight of the compound is twice that of the empirical formula. Therefore the molecular formula of your compound is:2 *(NO2) ---> N2O4

it can be calculated using the formula percentage composition of N =Gram molecular mass of nitrogen in the compound/ Gram molecular mass of compound *100

On the periodic table you will find the two elements in a binary molecular compound on the right side. Every binary molecular compound name contains a prefix.

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