How to Calculate Molecular Weight (Worked Examples)

Molecular weight is the mass of one molecule of a substance, found by adding up the atomic weights of every atom in its formula. That is the whole idea. The tricky parts are reading the formula correctly (brackets, hydrates, subscripts), knowing which atomic weights to use, and turning the answer into something useful, such as how many grams to weigh out for a solution. This guide walks through all of it with fully worked examples.

Molecular weight, molar mass and Mr: are they the same?

You will see several names for nearly the same number, and it helps to know which one your course or lab uses:

• Relative formula mass (M_r) is the term used in UK GCSE and A-level chemistry. It has no units because it is a ratio compared with one-twelfth of a carbon-12 atom. Exam boards give you rounded relative atomic masses (A_r), such as H = 1, C = 12, O = 16 and Cl = 35.5.
• Molar mass is the term most US high school and college courses prefer. It is the mass of one mole of the substance and is written in grams per mole (g/mol).
• Molecular weight is the older, everyday name, still printed on most chemical bottles and supplier catalogues. Numerically it matches the molar mass.
• Molecular mass in daltons (Da) is common in biochemistry. One dalton is one-twelfth of the mass of a carbon-12 atom, so a protein of 50,000 Da has a molar mass of about 50,000 g/mol.

For water, all four give the same figure: about 18. The only real differences are the units and how many decimal places you keep.

Which atomic weights should you use?

Atomic weights come from the periodic table. They are rarely whole numbers because most elements are a mix of isotopes. Chlorine, for example, is roughly 75.8% chlorine-35 (mass 34.969) and 24.2% chlorine-37 (mass 36.966). The weighted average is:

0.758 × 34.969 + 0.242 × 36.966 = 26.506 + 8.946 = 35.452

which is why chlorine is listed as 35.45 rather than 35 or 37. IUPAC publishes the official standard atomic weights. For everyday work, these values are more than precise enough:

Element	Symbol	Atomic weight (lab use)	GCSE/exam value
Hydrogen	H	1.008	1
Carbon	C	12.011	12
Nitrogen	N	14.007	14
Oxygen	O	15.999	16
Sodium	Na	22.990	23
Sulfur	S	32.06	32
Chlorine	Cl	35.45	35.5
Calcium	Ca	40.078	40
Copper	Cu	63.546	63.5

Rule of thumb: in an exam, use exactly the values on the data sheet you are given, because the mark scheme is built on them. In a lab, use the values to at least two or three decimal places. You can look any element up on our interactive periodic table.

How to calculate molecular weight: step by step

1. Write the formula clearly. Check capital letters: Co is cobalt, CO is carbon monoxide.
2. Count the atoms of each element. A subscript applies only to the symbol directly before it. A subscript after a bracket multiplies everything inside the bracket.
3. Look up each atomic weight.
4. Multiply each atomic weight by the number of those atoms.
5. Add the results together. Give the answer in g/mol (or with no units if you are quoting M_r).

Worked examples, line by line

Example 1: Water, H₂O

Atoms: 2 hydrogen, 1 oxygen.

Hydrogen: 2 × 1.008 = 2.016
Oxygen: 1 × 15.999 = 15.999
Total: 2.016 + 15.999 = 18.015 g/mol

Example 2: Glucose, C₆H₁₂O₆

Carbon: 6 × 12.011 = 72.066
Hydrogen: 12 × 1.008 = 12.096
Oxygen: 6 × 15.999 = 95.994
Total: 72.066 + 12.096 + 95.994 = 180.156 g/mol

With GCSE values the same sum is 72 + 12 + 96 = 180, which is the answer a UK mark scheme expects.

Example 3: Brackets, ammonium sulfate (NH₄)₂SO₄

The 2 after the bracket doubles everything inside it, so there are 2 nitrogen and 8 hydrogen atoms, plus 1 sulfur and 4 oxygen.

Nitrogen: 2 × 14.007 = 28.014
Hydrogen: 8 × 1.008 = 8.064
Sulfur: 1 × 32.06 = 32.06
Oxygen: 4 × 15.999 = 63.996
Total: 28.014 + 8.064 + 32.06 + 63.996 = 132.134 g/mol

The most common mistake here is counting 4 hydrogens instead of 8. If your answer for a bracketed formula looks low, recount the bracket first.

Example 4: Hydrates, copper(II) sulfate pentahydrate CuSO₄·5H₂O

The dot means five water molecules are locked into each unit of the crystal, and they count towards the weight. Work out each part separately, then add.

CuSO₄: 63.546 + 32.06 + (4 × 15.999) = 63.546 + 32.06 + 63.996 = 159.602
5H₂O: 5 × 18.015 = 90.075
Total: 159.602 + 90.075 = 249.677 g/mol

This matters in practice. The blue crystals on a school shelf are the pentahydrate, and if you use 159.6 instead of 249.7 in a calculation, you will weigh out far too little copper sulfate.

Putting it to work: grams, moles and solutions

Molecular weight is the bridge between the mass you can weigh and the number of moles a reaction cares about:

moles = mass (g) ÷ molar mass (g/mol)
mass (g) = moles × molar mass (g/mol)

Worked example: you need 250 mL of 0.5 mol/L sodium chloride solution.

Molar mass of NaCl: 22.990 + 35.45 = 58.44 g/mol
Volume in litres: 250 mL = 0.250 L
Moles needed: 0.5 × 0.250 = 0.125 mol
Mass needed: 0.125 × 58.44 = 7.305 g (about 7.3 g)

For readers working in imperial units: 7.305 g is about 0.26 oz (7.305 ÷ 28.35), and 250 mL is about 8.45 US fluid ounces or 8.80 UK fluid ounces. Laboratory work is done in metric almost everywhere, so it is worth converting once and then staying in grams and millilitres. If you then need to dilute that stock solution, the solution dilution calculator handles the C₁V₁ = C₂V₂ step.

You can also use molecular weight to find percentage composition. For water, oxygen makes up 15.999 ÷ 18.015 = 0.8881, or about 88.8% of the mass, and hydrogen the remaining 11.2%.

Frequently asked questions

Is molecular weight the same as molar mass?

Numerically, yes. Molar mass is expressed in g/mol, while molecular weight (or M_r) is technically a unitless ratio. Water is 18.015 either way, so you can use them interchangeably in calculations as long as you label the units your teacher or journal expects.

Why does my answer differ slightly from the one on the bottle?

Different sources round atomic weights differently, and some older tables use values such as S = 32.065. A difference in the second decimal place is normal. A difference of several whole units usually means a missed bracket, a missed hydrate, or the wrong salt form.

How many decimal places should I give?

For UK GCSE, follow the data sheet, which usually means whole numbers or one decimal place. For lab work and A-level or college chemistry, two decimal places (for example 58.44 g/mol for NaCl) is the usual standard.

Check your working in seconds

Doing the sum by hand is the best way to understand it, but for long formulas it is easy to slip. Type your formula into the free Quoteen molecular weight calculator to check the molar mass against your own working. For other chemistry sums, try the chemistry calculator.