Dissolution Thermochemistry

Systematically measure which salts heat or cool water — and calculate why

Difficulty: Easy | Time: 30–45 min | Visual Impact: Medium

Historical Context

The distinction between exothermic and endothermic processes was cemented in the mid-19th century as thermodynamics matured. Hess’s law (1840) established that energy changes in chemical reactions are additive regardless of path, which made it possible to tabulate dissolution energies systematically. Scientists noticed that some salts heated their containers while others cooled them — a puzzle whose answer turned out to hinge on how strongly each ion grips water molecules.

Commercial instant cold packs were developed in the 1960s using ammonium nitrate (now often replaced with ammonium chloride or urea due to safety concerns). Hot packs using calcium chloride followed. These products brought thermochemistry into everyday life — and into sports medicine bags worldwide.

This experiment takes the commercial pack demonstration further: instead of just observing hot versus cold, you will test seven different solutes, rank them, and use simple calorimetry to estimate their heats of dissolution from your own measurements.

Materials

Solute Amount From
Calcium chloride (CaCl₂) 15 g Chemical inventory
Ammonium chloride (NH₄Cl) 15 g Chemical inventory
Sodium chloride (table salt, NaCl) 15 g Kitchen
Sodium carbonate (Na₂CO₃) 15 g Chemical inventory
Epsom salt (MgSO₄·7H₂O) 15 g Chemical inventory
Baking soda (NaHCO₃) 15 g Kitchen
Citric acid 15 g Chemical inventory

Equipment: - Thermometer (ideally digital, reading to 0.1°C) - 7 × 250 mL cups or beakers - 7 × 100 mL portions of water, all at the same starting temperature - Kitchen scale or measuring spoons - Stirring rod or spoon - Notebook for recording temperatures

Procedure

Setting up

  1. Measure 100 mL of water into each of seven cups. Let them all stand together for 5 minutes so they reach the same temperature — small differences in starting temperature will throw off the comparison.
  2. Record the starting temperature of each cup (should be the same, within 0.5°C).
  3. For each solute, weigh out exactly 15 g.

Dissolving and measuring

  1. Add the first solute (calcium chloride) to cup 1, all at once, and stir vigorously for 30 seconds.
  2. Record the temperature at 30 s, 60 s, and 2 min. Note the maximum (or minimum) temperature reached.
  3. Repeat for all seven solutes, working quickly so the water temperature does not drift.

Take care with calcium chloride — it dissolves vigorously and can splash. Add it gradually if needed.

Recording results

Solute Starting T (°C) Final T (°C) ΔT (°C) Hot or cold?
CaCl₂
Na₂CO₃
NaCl
MgSO₄·7H₂O
NH₄Cl
NaHCO₃
Citric acid

Typical results you should see: CaCl₂ heats to about 45–55°C; NH₄Cl and NaHCO₃ cool noticeably; citric acid cools moderately; NaCl barely changes.

Simple Calorimetry

Once you have your temperature changes, you can estimate the heat of dissolution for each solute — the energy released or absorbed per mole.

Step 1 — Calculate energy transferred to the water:

\[q = m \times c \times \Delta T\]

Where: - \(m\) = mass of water = 100 g - \(c\) = specific heat of water = 4.18 J g⁻¹ °C⁻¹ - \(\Delta T\) = temperature change in °C

For example, if calcium chloride raised the temperature by 30°C:

\[q = 100 \times 4.18 \times 30 = 12{,}540 \text{ J} = 12.5 \text{ kJ}\]

The water gained 12.5 kJ of heat, so the reaction released 12.5 kJ (the reaction’s \(q\) is −12.5 kJ).

Step 2 — Convert to kJ per mole:

\[\Delta H_{\text{measured}} = \frac{-q}{n}\]

For CaCl₂ (molar mass = 111 g/mol), 15 g = 0.135 mol:

\[\Delta H = \frac{-12.5 \text{ kJ}}{0.135 \text{ mol}} = -93 \text{ kJ/mol}\]

Compare this with other solutes using their molar masses: - NH₄Cl: 53.5 g/mol - NaCl: 58.4 g/mol - Na₂CO₃: 106 g/mol - MgSO₄·7H₂O: 246 g/mol - NaHCO₃: 84 g/mol - Citric acid: 192 g/mol

Note on accuracy: Your measurements will differ from literature values because heat also goes into warming the cup, the thermometer, and the surrounding air. Student calorimetry typically captures 50–80% of the true heat. The direction (exo vs. endo) and relative ranking will be accurate even if the absolute numbers are not.

Literature Values

How do your results compare?

Solute Literature ΔH (kJ/mol) Typical ΔT with 15 g in 100 mL
CaCl₂ −81 +30 to +40°C
Na₂CO₃ −28 +8 to +14°C
NaCl +3.9 −1 to +1°C
MgSO₄·7H₂O +13 −3 to −6°C
NH₄Cl +15 −5 to −8°C
NaHCO₃ +27 −6 to −9°C
Citric acid +23 −5 to −8°C

Notice: the ranking from hottest to coldest spans nearly 110 kJ/mol — a huge range for such simple processes.

The Science

Why does dissolving produce heat at all?

Dissolution always involves two competing energy changes:

1. Breaking the crystal lattice — always endothermic

Ions in a solid are held in place by strong electrostatic attraction. Pulling them apart requires energy input. The strength of this attraction is the lattice energy, which scales with ion charge and decreases with ion size.

2. Hydrating the ions — always exothermic

Once free in solution, each ion is surrounded by water molecules oriented with their partial charges facing the ion. This hydration releases energy. Small ions with high charge attract water more strongly, releasing more energy on hydration.

Net effect: If hydration releases more energy than lattice-breaking requires → exothermic (solution heats). If lattice-breaking costs more than hydration releases → endothermic (solution cools).

Why is CaCl₂ so hot?

Calcium(II) is a small, doubly charged ion. Its high charge density pulls water molecules extremely tightly, releasing a large hydration energy that more than compensates for the lattice energy. The result is a strongly exothermic dissolution. Sodium carbonate is moderately exothermic for similar reasons — Na⁺ and CO₃²⁻ both hydrate well.

Why is NaHCO₃ cold?

Bicarbonate (HCO₃⁻) is a large, singly charged ion. It does not pull water as tightly, so hydration energy is modest. But the lattice energy is still significant, leaving a net endothermic balance.

Why does MgSO₄·7H₂O cool water even though anhydrous MgSO₄ would heat it?

This is the subtlest result. Anhydrous magnesium sulfate (MgSO₄) has a strongly exothermic dissolution (about −91 kJ/mol) — Mg²⁺ is small and very highly charged, so it hydrates powerfully. But Epsom salt is already hydrated: each formula unit comes with 7 water molecules attached. Before Mg²⁺ can hydrate from bulk water, it must first shed those seven water molecules — an endothermic step that reverses the advantage. The net result tips slightly endothermic.

This is the reason anhydrous CaCl₂ (sometimes sold as a desiccant) heats more dramatically than partially-hydrated forms. The more dehydrated the salt, the more exothermic the dissolution.

Charge density summary

Ion Charge Radius Hydration enthalpy (kJ/mol)
Ca²⁺ 2+ small −1592
Mg²⁺ 2+ very small −1922
Na⁺ 1+ medium −406
NH₄⁺ 1+ large −301
Cl⁻ 1− medium −363
SO₄²⁻ 2− large −1138
HCO₃⁻ 1− large −335

High charge and small size = strong hydration = more exothermic dissolution (all else equal).

Explore Further

Dehydration reversal: Pour your hot calcium chloride solution into a flat dish and leave it in a warm place to evaporate slowly. As the water evaporates, CaCl₂ crystals reform. Touch the dish near the end — is it warm? The reverse process (crystallization) releases the same energy as dissolution absorbed; it is exothermic.

The body heat connection: Chemical hand warmers use a pouch of supersaturated sodium acetate or iron oxidation — both exothermic processes. How does their sustained heat compare to the brief spike you observed with CaCl₂? (The sodium acetate Hot Ice experiment explores this.)

Measuring heat capacity: If you add the same mass of ethanol instead of water, the temperature change will be different because ethanol has a lower specific heat (2.44 J/g·°C vs. 4.18 for water). Try dissolving a small amount of NaCl in ethanol and compare the ΔT to the water result — then recalculate: does the calculated ΔH come out the same?

Chemicals used in this experiment: