Crystalloluminescence of Sodium Chloride
| Polish version is here |
The following article was originally published in the journal for educators Fizyka w szkole (eng. Physisc in School) (3/2014):
A crystal, also called a crystalline solid, is a solid material in which atoms, molecules, or ions are arranged in a highly ordered microscopic structure, forming a crystal lattice that extends in all directions.
Because of their structure, crystals are often seen as a symbol of geometric beauty. Escher, the renowned Dutch painter and graphic artist, emphasized this multiple times. He wrote that “long before there were people on the earth, crystals were already growing in the earth's crust. On one day or another, a human being first came across such a sparkling morsel of regularity lying on the ground or hit one with his stone tool and it broke off and fell at his feet, and he picked it up and regarded it in his open hand, and he was amazed” [1].
Crystalline substances certainly inspired many of his works. This is apparent in the 1949 woodcut entitled “Double Planetoid” (Fig. 1).
Crystals exhibit many intriguing properties, such as piezoelectricity or thermoluminescence.
Another fascinating phenomenon is crystalloluminescence — luminescence produced during crystallization. It can even be observed in a substance as common as sodium chloride.
A challenge, however, is that experiments described in older books often fail to produce the desired result. Although this phenomenon was first described by Ernest Bandrowski back in 1894, it was not until almost 120 years later that Andrew J. Alexander outlined the conditions for more reproducible observations. The experiment described below is based on the procedure he described [2].
Preparations
To carry out the experiment, you’ll need the following substances:
- hydrochloric acid HCl 24%
- sodium chloride NaCl
- silver nitrate AgNO3
- copper(II) sulfate pentahydrate CuSO4·5H2O
Warning: Hydrochloric acid HCl is corrosive. The hydrogen chloride is irritating, and at higher concentrations can be toxic. Silver nitrate and copper(II) sulfate, like most heavy-metal salts, are toxic. Avoid any skin contact with silver salts — exposure to light will create difficult-to-remove black stains.
The easiest substance to obtain is sodium chloride NaCl, which is used in nearly pure form as table salt. This chloride forms beautiful cubic crystals (Photo. 1).
Commercial salt often contains a tiny amount of iodine compounds to protect consumers from conditions such as goiter. It is best to use the cheapest, non-iodized salt.
When preparing all solutions, make sure to use distilled water! You must strictly adhere to the stated concentrations. For the experiment to succeed, all glassware must be kept very clean.
First, prepare the required solutions. To make solution A, dissolve enough sodium chloride in 100 cm3 (3.38 fl oz) of water to form a saturated solution at room temperature, then add 5 cm3 (0.17 fl oz) of a 0.07M silver nitrate solution. A white precipitate of silver chloride appears in the process (Photo. 2).
Although silver chloride AgCl is practically insoluble in pure water, it dissolves relatively well in concentrated sodium chloride solutions, forming chloride complexes such as AgCl2-, AgCl32-, and AgCl43-. You can confirm this by gently shaking the solution — after a moment, the precipitate disappears and the liquid becomes clear (Photo. 3).
To make solution B, add 0.5 cm3 (0.017 fl oz) of a 0.05M copper(II) sulfate pentahydrate solution to 100 cm3 (3.38 fl oz) of 24% hydrochloric acid. The solution takes on a greenish hue (Photo. 4), caused by the formation of copper(II) chloride complexes.
After preparing these solutions, cover them and let them stand for several hours so that any solid impurities have time to settle.
The Experiment
For the experiment, measure equal volumes of solutions A and B — e.g., 10–15 cm3 (0.34–0.51 fl oz) — into separate small beakers. The solutions must be clear — the presence of even tiny crystals of undissolved substances or dust can prevent the desired effect! Next, darken the room and allow your eyes to adjust. Gently pour solution A into the beaker containing solution B. The way you pour and mix is crucial, and you may need some practice before you achieve satisfactory results. The experiment can be somewhat unpredictable — if it doesn’t work the first time, don’t be discouraged; just try again.
If everything is set up correctly, you’ll notice faint bluish sparks after a few seconds, especially near the bottom of the container. The light flashes are very brief but are relatively easy to see. The entire phenomenon usually lasts about a minute.
This crystalloluminescence can be recorded using a long exposure photo (Photo. 5; ISO400, 60s).
You can see that the crystalloluminescence indeed appears as random, isolated bursts of light throughout the liquid volume. Each flash of light is caused by the formation of a tiny sodium chloride crystal. With the light on, you can see these crystals slowly settling at the bottom of the beaker (Photo. 6).
Explanation
When hydrochloric acid is mixed with a saturated sodium chloride solution, the solubility of sodium chloride decreases. This is due to an increased concentration of chloride ions Cl−. Sodium chloride then precipitates from the solution as crystals.
If there are no crystallization nuclei in the solution — such as undissolved salt grains, air bubbles, or dust particles — the crystallization process leads to the sudden formation of many small crystals.
The energy from stresses within the forming crystal lattice is released partly as electromagnetic radiation — blue and ultraviolet light. Crystallographic defects promote this, so we deliberately introduce minor impurities, such as Cu2+ or Ag+ ions, into the lattice. In pure sodium chloride (NaCl), this effect would be imperceptible or significantly weaker and harder to detect.
References:
- [1] Hodorowicz S. A., Krótka refleksja o krysztale i krystalografii, Alma Mater nr 136, 2011, p. 28
- [2] Alexander A. J., Deep ultraviolet and visible crystalloluminescence of sodium chloride, Journal of chemical physics 2012, vol. 136
All photographs and illustrations were created by the author.
Addendum
In the published article, Photo 5 was presented with the container’s outline marked for improved clarity. Below is the unedited version of the image:
Marek Ples