How to Name Hydrates

Blue copper(II) sulfate crystals turn into a white powder when heated, and back to blue when water is added. The blue crystals aren’t pure copper(II) sulfate: they contain water molecules locked into the crystal structure. Compounds like this are called hydrates, and they have their own naming rules and formula notation. This guide explains how to name them, how to write their formulas, and how to work out the amount of water from an experiment.

What is a hydrate?

A hydrate is a compound — usually an ionic salt — whose crystals contain a fixed number of water molecules per formula unit. This water is called water of crystallisation (see hydrates and water of crystallisation).

The water molecules aren’t just damp water on the surface. They’re part of the crystal structure, often bonded to the metal ion through dative bonds from water’s oxygen lone pairs (see dative covalent bonds), or held by hydrogen bonds to the anions. In blue copper(II) sulfate, four water molecules surround each Cu²⁺ ion, and the fifth is hydrogen-bonded between sulfate ions.

When a hydrate loses its water — usually by heating — the product is called anhydrous (“without water”).

Writing hydrate formulas: the dot

A hydrate’s formula is written as the salt formula, then a raised dot (·), then the number of water molecules:

CuSO₄·5H₂O

The dot means “combined with”. It doesn’t mean multiply. So CuSO₄·5H₂O means one CuSO₄ unit combined with five H₂O molecules.

(If you can’t type a raised dot, a full stop or an asterisk is often accepted, e.g. CuSO₄.5H₂O.)

Naming hydrates: Greek prefixes

A hydrate’s name has two parts:

1. The name of the salt (following the normal rules for ionic compounds — see how to name ionic compounds).
2. A Greek prefix showing the number of water molecules, followed by “hydrate”.

Number of H₂O	Prefix	Word
1	mono-	monohydrate
2	di-	dihydrate
3	tri-	trihydrate
4	tetra-	tetrahydrate
5	penta-	pentahydrate
6	hexa-	hexahydrate
7	hepta-	heptahydrate
8	octa-	octahydrate
9	nona-	nonahydrate
10	deca-	decahydrate
½	hemi-	hemihydrate

Step-by-step: name from formula

1. Name the ionic compound in front of the dot. Include a Roman numeral for transition metals with more than one possible charge (see naming compounds with transition metals).
2. Count the water molecules after the dot.
3. Add the Greek prefix + “hydrate”.

Worked examples

Formula	Name
CuSO₄·5H₂O	copper(II) sulfate pentahydrate
CaCl₂·2H₂O	calcium chloride dihydrate
MgSO₄·7H₂O	magnesium sulfate heptahydrate (Epsom salt)
Na₂CO₃·10H₂O	sodium carbonate decahydrate (washing soda)
CoCl₂·6H₂O	cobalt(II) chloride hexahydrate
FeSO₄·7H₂O	iron(II) sulfate heptahydrate
CaSO₄·2H₂O	calcium sulfate dihydrate (gypsum)
CaSO₄·½H₂O	calcium sulfate hemihydrate (plaster of Paris)
Na₂SO₄·10H₂O	sodium sulfate decahydrate (Glauber’s salt)

Step-by-step: formula from name

1. Write the formula of the salt from its name, balancing charges.
2. Convert the prefix into a number.
3. Write the dot and the water.

Example: nickel(II) chloride hexahydrate

• Ni²⁺ and Cl⁻ → NiCl₂
• hexa- = 6
• NiCl₂·6H₂O

Example: sodium thiosulfate pentahydrate

• 2Na⁺ and S₂O₃²⁻ → Na₂S₂O₃
• penta- = 5
• Na₂S₂O₃·5H₂O

Anhydrous forms

The anhydrous salt is named without any hydrate part, sometimes with the word “anhydrous” added for clarity:

• CuSO₄ — anhydrous copper(II) sulfate (white powder)
• CuSO₄·5H₂O — copper(II) sulfate pentahydrate (blue crystals)

Some hydrates lose water in stages. Gypsum (CaSO₄·2H₂O), heated to about 150 °C, becomes the hemihydrate (plaster of Paris). When water is mixed back in, it re-forms gypsum and sets hard — the chemistry behind plaster casts and plasterboard.

Calculating molar mass

When calculating the molar mass of a hydrate, include the water:

CuSO₄·5H₂O:

• CuSO₄ = 63.55 + 32.06 + 4(16.00) = 159.61 g mol⁻¹
• 5H₂O = 5 × 18.02 = 90.10 g mol⁻¹
• Total = 249.71 g mol⁻¹

The molar mass calculator accepts hydrate formulas such as CuSO4·5H2O and shows the breakdown, including the mass percentage of water.

Percentage of water in CuSO₄·5H₂O = 90.10 ÷ 249.71 × 100 = 36.1 %

This matters in the lab: if you weigh out 10 g of blue copper(II) sulfate, only about 6.4 g of it is actually CuSO₄.

Finding the formula of a hydrate by experiment

A classic practical: heat a weighed sample of a hydrate to drive off the water, and weigh the anhydrous residue.

Worked example: A student heats 5.00 g of hydrated magnesium sulfate, MgSO₄·xH₂O, until the mass is constant. The anhydrous MgSO₄ left has a mass of 2.44 g. Find x.

1. Mass of water lost = 5.00 − 2.44 = 2.56 g
2. Moles of MgSO₄ = 2.44 ÷ 120.37 = 0.02027 mol (M of MgSO₄ = 24.31 + 32.06 + 64.00 = 120.37 g mol⁻¹)
3. Moles of H₂O = 2.56 ÷ 18.02 = 0.1421 mol
4. Ratio H₂O : MgSO₄ = 0.1421 ÷ 0.02027 = 7.01
5. x = 7, so the formula is MgSO₄·7H₂O, magnesium sulfate heptahydrate.

Good practice in this experiment:

• Heat to constant mass — reheat and reweigh until two readings agree — to make sure all the water has gone.
• Avoid overheating, which can decompose some salts further (for example, releasing SO₃ from sulfates).
• Cool in a desiccator so the anhydrous salt doesn’t reabsorb water from the air before weighing.

Hydrates you might meet

• Cobalt(II) chloride paper is blue when dry (anhydrous or lower hydrate) and pink when hydrated — a simple test for water.
• Anhydrous copper(II) sulfate turns from white to blue with water — another test for water.
• Silica gel sachets in packaging absorb moisture; some contain a coloured indicator that changes as it becomes hydrated.

Why hydrates matter in the lab

Hydrates catch students out in practical work more often than in written exams. If a method says “dissolve 2.50 g of copper(II) sulfate”, it matters whether the bottle contains the anhydrous white powder or the blue pentahydrate. The same mass of the pentahydrate contains only about 64 % as much copper(II) sulfate, so a solution made from it would be much more dilute than intended. Good practice is to read the label carefully, check the formula given, and use the correct molar mass in every calculation.

Some anhydrous salts are also useful precisely because they’re so eager to become hydrates. Anhydrous calcium chloride and anhydrous magnesium sulfate are used as drying agents: they pull water out of damp organic liquids or gases by forming hydrates. And some hydrates slowly lose water to the air (they effloresce), while other substances absorb water from the air until they dissolve (they’re deliquescent), so stored samples can change composition over time.

Common mistakes

• Treating the dot as multiplication. CuSO₄·5H₂O is not 5 × CuSO₄ × H₂O.
• Forgetting the water when calculating molar mass.
• Using the wrong prefix — hepta is 7, hexa is 6, octa is 8.
• Leaving out Roman numerals for transition metals — “copper sulfate pentahydrate” should be copper(II) sulfate pentahydrate.

Key takeaways

• A hydrate contains a fixed number of water molecules per formula unit, shown with a dot: CuSO₄·5H₂O.
• Name = salt name + Greek prefix + “hydrate” (e.g. copper(II) sulfate pentahydrate).
• The water-free compound is anhydrous.
• Always include water in molar mass calculations.
• Find x experimentally by heating to constant mass and comparing moles of water to moles of anhydrous salt.