Empirical Formula

Empirical formula

Slide 2

• Empirical Formula - A formula that gives the simplest whole-number ratio of atoms in a compound.
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Slide 3

• Start with the number of grams of each element, given in the problem.
• If percentages are given, assume that the total mass is 100 grams so that
• the mass of each element = the percent given.
• Convert the mass of each element to moles using the molar mass from the periodic table.
• Divide each mole value by the smallest number of moles calculated.
• Round to the nearest whole number. This is the mole ratio of the elements and is
• represented by subscripts in the empirical formula.
• If the number is too far to round (x.1 ~ x.9), then multiply each solution by the same
• factor to get the lowest whole number multiple.
• e.g. If one solution is 1.5, then multiply each solution in the problem by 2 to get 3.
• e.g. If one solution is 1.25, then multiply each solution in the problem by 4 to get 5.
• Once the empirical formula is found, the molecular formula for a compound can be determined if the molar mass of the compound is known. Simply calculate the mass of the empirical formula and divide the molar mass of the compound by the mass of the empirical formula to find the ratio between the molecular formula and the empirical formula. Multiply all the atoms (subscripts) by this ratio to find the molecular formula.

My most common mistake

• My most common mistake is the adding of the number at the end of the letters to both of the letters in the formula.
• The only time that you could use the little number with both of the letters is when they are in brackets e.g. (NH)3+O2 -- NO+H2O this would allow you to use the little nuimber
• do not use this number with both of the letters unless they are in brackets ()
• NH3+O2 ---- NO+H2O