Molecular Mass Notes

I. Molar mass: The mass of one atom in atomic mass units is the atomic mass from the periodic table, but the mass of one atom in grams is infinitesimal. However, if we take a very large number of atoms we can weigh them in grams. If we take 6.02 x 1023 atoms (1 mole), their weight in grams will be equal to the weight of one in amu.

1 mole of X is just 6.02 x 1023 of X. X can be atoms, molecules, ions, formula units, donuts, or anything else we want to count.

If we take a mole of molecules or ionic compounds and weigh it we will get the molar mass, which is the same numeric value as the formula mass, but will be measured in grams.

II. Conversions Our conversion factors: 1 mole = 6.022 x 1023 atoms 1 mole = (molar mass)grams We can use a T chart system to convert between atoms, moles and grams. Problems will always be set up like: GIVEN INFO. WITH UNITS # UNITS ASKED TO CONVERT INTO ______

# UNITS OF THE GIVEN INFORMATION

The numbers always come from the conversion factors

III. Formula or molecular mass To find the formula mass of an ionic compound or the molecular mass of a covalently bonded compound simply multiply the average atomic mass of each constituent element by the number of that element present in the compound, and then add each of those products together.

IV. Percent Composition The percent composition of a compound is the percentage of the total mass of a compound contributed by the constituent parts. To find: 1-Find the mass of compound. 2-Divide the total mass of each constituent part by the mass of the entire compound. Multiply by 100.

V. Empirical Formula An empirical formula represents the simplest whole number ratio of the atoms in a compound. If we know the opposite charges on the particles making up a compound, we can find this ratio by the LCM or criss cross method. For ionic compounds with transition metals if we are not given a roman numeral or a formula we have to find the empirical formula experimentally. To do this: 1) Convert all gram or percent measurements to moles 2) Divide all mole amounts by the smallest mole amount. 3) Multiply to get rid of fractions if necessary. 4) Use your answers as your subscripts.

VI. Molecular Formula Molecular Formula--Molecules do not always have the simplest whole number ratio of atoms, so we write the actual ratio of atoms in a molecule. This must be calculated experimentally. To do this : 1--Find the empirical formula. 2--Find the molecular mass of the empirical formula. 3--Divide the actual molecular mass by the empirical formula's molecular mass. 4--Multiply the subscripts in the empirical formula by the answer from part three.

VII. Empirical Formula of a Hydrate Hydrates are compounds that have water molecules trapped within their crystalline structure…think of a sponge…the water is not glued to the sponge it just gets trapped in the little holes of the sponge…water is not chemically bonded to the hydrate just trapped in the gaps in its structure. Different amounts of water can get stuck in different compounds because they have different sized holes. To find their formula: 1—Find the mass of the hydrate 2—Dry the hydrate turning it into an “anhydrate” 3—Subtract the two masses to find the mass of water removed 4—Divide the mass of the anhydrate by the molar mass of the anhydrate to get the moles of anhydrate 5—Divide the mass of water removed by the molar mass of water to get the moles of water 6-- Divide both mole amounts by the smallest mole amount to get the ratio of anhydrate to water X : Y 7—Write the formula X anhydrate * Y H2O