How to Calculate Molar Mass From a Chemical Formula (With Worked Examples)

Molar mass is the bridge between the formula on a page and the grams you weigh on a balance. Once you can read a chemical formula and turn it into a single number in g/mol, the rest of intro chemistry — making a solution, balancing a reaction by mass, checking a percent composition — stops feeling like guesswork. This guide walks through the one rule that does all the work, then shows how parentheses and hydrates fit the same rule, and ends with the mole-to-gram conversion you actually use in lab.

The one rule: atomic weight times subscript, summed

Here is the whole method in a sentence: molar mass equals the sum, over every element in the formula, of that element's atomic weight multiplied by how many atoms of it the formula contains. That subscript count is the number written after the symbol — and an absent subscript means one.

So the procedure is mechanical:

List each distinct element in the formula.
Count its atoms (the subscript, defaulting to 1).
Multiply that count by the element's standard atomic weight.
Add up all the contributions.

The atomic weights come straight off the periodic table. I lean on the 2021 IUPAC conventional values — hydrogen 1.008, carbon 12.011, oxygen 15.999, and so on — because those are the numbers printed on a modern chart, so your answer lands on the textbook figure instead of drifting in the last digit. If you want the full table of weights in front of you while you work, the periodic table reference has every element's value.

Worked example: water and sulfuric acid

Start with water, H₂O. Two elements, two contributions:

Hydrogen: 2 atoms × 1.008 = 2.016
Oxygen: 1 atom × 16.00 = 16.00

Sum: 18.02 g/mol. That is the whole calculation. The 18 you half-remember is just 2×1.008 + 15.999 rounded to four significant figures.

Now a meatier one — sulfuric acid, H₂SO₄:

Hydrogen: 2 × 1.008 = 2.016
Sulfur: 1 × 32.06 = 32.06
Oxygen: 4 × 15.999 = 63.996

Sum: 98.08 g/mol. Nothing new happened. The oxygen subscript is 4 instead of 1, so you multiply by 4 before adding. Every formula you will ever meet is this same loop, just with more rows.

A quick sanity check you can always run: each element's contribution divided by the total is its mass percent, and those percentages have to add to 100. For water, oxygen is 16.00 / 18.02 ≈ 88.8% and hydrogen is the remaining 11.2%. If your percentages don't sum to 100, you miscounted an atom somewhere.

Parentheses and brackets: expand, then sum

Subscripts on a parenthesis multiply everything inside the group. Calcium hydroxide, Ca(OH)₂, has that subscript 2 hanging off the closing parenthesis, so it doubles the whole OH group:

Calcium: 1 × 40.078 = 40.078
Oxygen: 2 × 15.999 = 31.998 (the O inside, taken twice)
Hydrogen: 2 × 1.008 = 2.016 (the H inside, taken twice)

Sum: 74.09 g/mol.

Nested brackets follow the same logic from the inside out. In potassium ferrocyanide, K₄[Fe(CN)₆], you resolve the inner (CN)₆ first — 6 carbons and 6 nitrogens — then apply the outer bracket, then add the 4 potassiums and 1 iron. Work the innermost group first and the rest falls into place; the total comes out to 368.35 g/mol.

The trap here is ammonium sulfate, (NH₄)₂SO₄. The 2 multiplies both the nitrogen and the four hydrogens inside the parentheses, giving 2 nitrogens and 8 hydrogens, not 1 and 4. Distribute the group multiplier across every atom inside before you add the sulfur and oxygen.

Hydrates: the dot adds water

A hydrate dot — written ·, •, or just a period — separates the formula into parts, and the number right after the dot multiplies that entire following part. Copper(II) sulfate pentahydrate, CuSO₄·5H₂O, is anhydrous CuSO₄ plus five whole water molecules:

From CuSO₄: 1 Cu, 1 S, 4 O
From 5H₂O: 10 H, 5 O

So the oxygen count is 4 + 5 = 9, and the molar mass is 249.69 g/mol — versus 159.61 for the anhydrous salt. That difference is the part students most often drop, and it is not small: weigh out the anhydrous mass when your bottle is the pentahydrate and your solution is off by more than a third. Always match the hydration state on the label.

From molar mass to grams: n = m / M

Molar mass earns its keep the moment you need to weigh something. The relationship is n = m / M — moles equal mass divided by molar mass — and rearranged, m = n × M gives you grams from moles.

Say a protocol calls for 0.25 mol of sodium chloride. NaCl has molar mass 58.44 g/mol, so:

m = 0.25 mol × 58.44 g/mol = 14.61 g

That is the number you dial into the balance. No scribbling the formula on a paper towel, no transposing a digit between the calculation and the weighing boat. When I'm setting up a bench prep, I run the formula through the molar mass calculator, flip the built-in converter to moles → mass, and copy the gram figure straight onto my prep sheet — it has caught me more than once when I'd misremembered a hydrate's weight and would otherwise have weighed out the anhydrous mass. Once you have the gram figure and a target volume, a dilution calculator handles the C₁V₁ = C₂V₂ step for getting to your final concentration.

Common slip-ups worth naming

A few mistakes show up over and over, and all of them produce a wrong number quietly rather than an error:

Case in element symbols. Co is cobalt; CO is carbon monoxide; Cu is copper, but CU is nothing. Capitalization is meaning, not style — a sloppy "NACL" is not the same as "NaCl".
Skipping the water of crystallization. Covered above, but it bears repeating because it is the single most expensive omission in solution prep.
Quoting too many decimals. Standard atomic weights carry real uncertainty, so reporting glucose as 180.156 g/mol claims a precision the underlying weights don't have. Four significant figures (180.2 g/mol) is honest.

Get the one rule — atomic weight times subscript, summed — and parentheses, brackets, and hydrates are all just careful bookkeeping on top of it. Count the atoms right, pull the weights off the table, add, and the grams follow.

Made by Toolora · Updated 2026-06-13