Article pubs.acs.org/jchemeduc
The Chemistry of Lightsticks: Demonstrations To Illustrate Chemical Processes Thomas Scott Kuntzleman,*,† Kristen Rohrer,† and Emeric Schultz‡ †
Department of Chemistry, Spring Arbor University, Spring Arbor, Michigan 49283, United States Department of Chemistry, Bloomsburg University, Bloomsburg, Pennsylvania 17815, United States
‡
S Supporting Information *
ABSTRACT: Lightsticks, or glowsticks as they are sometimes called, are perhaps the chemist’s quintessential toy. Because they are easy to activate and appealing to observe, experimenting with lightsticks provides a great way to get young people interested in science. Thus, we have used lightsticks to teach chemical concepts in a variety of outreach settings and demonstration shows. Although these devices are simple to operate, a working lightstick depends upon a rich array of physicochemical processes. For example, the chemical processes involved in lightsticks include acid−base chemistry, redox reactions, quantum chemistry, and thermodynamics. Consequently, we have used lightstick experiments and demonstrations in general, inorganic, and physical chemistry classes. In this paper, we share some experiments and demonstrations with lightsticks that we have used in these various educational settings. KEYWORDS: Elementary/Middle School Science, First-Year Undergraduate/General, General Public, High School/Introductory Chemistry, Upper-Division Undergraduate, Demonstrations, Acids/Bases, Dyes/Pigments, Spectroscopy, Thermodynamics
G
etting a lightstick to glow is simple: bend, and then shake well. Voila! There is an emission of light, the intensity and color of which depend upon the nature of the reaction inside. This simple device fascinates everyone who has ever experienced it. It is therefore no surprise that a variety of publications in this Journal and elsewhere1−7 describe the chemistry of lightsticks, the use of lightsticks in the laboratory, or the employment of lightsticks in chemical demonstrations. In this paper, we describe how the chemistry of lightsticks can provide an interesting and motivating backdrop around which a variety of topics in the undergraduate chemistry curriculum may be taught. Specifically, we describe lightstick experiments and demonstrations that can be performed to connect to a variety of concepts including acid−base chemistry, redox reactions, thermodynamics, and quantum chemistry. Because lightsticks are used as sources of chemical reagent(s) in the processes described, these experiments provide unique, facile, and inexpensive ways for educators to illustrate these chemical reactions. In this same spirit, we have made substantial efforts to devise these demonstrations so as to use familiar household materials or common chemicals found in most chemical stockrooms.
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For safety reasons, substituted phenyl oxalate esters are used so as to prevent formation of phenol; bis(2,4,6-trichlorophenyl)oxalate (TCPO) and bis(2,4,5-trichloro-6carbopentoxyphenyl)oxalate (CPPO) are the most commonly employed esters in lightstick formulations (Figure 1).
Figure 1. Structures of TCPO (left) and CPPO (right).
A lightstick consists of a flexible plastic casing with a central cavity. A glass vial in the cavity contains the oxalate, while the cavity around the vial contains the peroxide. A variety of solvents have been used in lightstick formulations, including alkyl phthalates, acetyl citrate esters, and alkyl benzoates.10 When the plastic outer casing is bent, the inner glass ampule breaks, the reactants mix, and the reaction begins. During the
BACKGROUND ON LIGHTSTICKS
Reactions in Lightsticks
The overall, light emitting reaction in lightsticks is the oxidation of phenyl oxalate esters by hydrogen peroxide to generate phenol derivatives and carbon dioxide.1,8,9 A base (often salicylate, C7H5O3−) is added as a catalyst: © 2012 American Chemical Society and Division of Chemical Education, Inc.
Published: May 15, 2012 910
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undergoes vibrational relaxation, VR (Figure 2, dashed lines), by transferring its vibrational energy to other molecules, typically to solvent through collisions. This radiationless transition does not emit light but has the effect of heating up the surrounding medium and bringing the dye molecule to its lowest vibrational level in S1. At this point, the dye molecule may emit a photon of light and relax to any one of several vibrational levels in S0 (Figure 2, CL: chemiluminescence). Alternatively, through a process called internal conversion, the dye could enter a very high-energy vibrational level of S0 that is at essentially the same energy as the ground vibrational level of S1 (Figure 2, IC: internal conversion). In this case, the dye molecule will release the excess vibrational energy through vibrational relaxation, increasing the temperature of the surroundings in the process. The chemiluminescence quantum yield,8,9 ϕCL, of the lightstick reaction is the number of photons of light (Nphoton) emitted for every molecule of reactant (Nreactant) consumed:
reaction between peroxide and the phenyl oxalate, a dioxetanedione intermediate (C2O4 in Scheme 1) is formed. Scheme 1. Mechanism Describing the Release of Energy in the Lightstick
When this highly reactive intermediate collides with the fluorescent dye molecule, energy is transferred to the fluorescent dye resulting in the promotion of an electron in the dye to an excited state.7−9 Concomitantly, C2O4 is cleaved into two molecules of CO2. Lastly, the dye molecule returns to its electronic ground state, releasing energy in the process. Much of this energy is released as light, but some is also released as heat. The emission of light from molecules by way of chemical excitation, such as this, is called chemiluminescence.
ϕCL =
Nphoton Nreactant
(2)
For a lightstick reaction, this value varies from 0.1% to 35% and depends upon a number of factors including the structure, reactivity, and concentrations of reactants and dyes.2,8,9,11 The chemiluminescent quantum yield is the product of the chemical yield of the chemiluminescent reaction (ϕc), the fraction of electronically excited-state dye molecules produced per reactant molecule (ϕe), and the fraction of electronically excited dye molecules that emit a photon per reactant molecule (ϕ′):
Energetics of Lightsticks
Whatever the quantity of energy released during the lightstick reaction, a portion of this energy (Figure 2, CE: chemical
ϕCL = ϕϕϕ ′ c e
(3)
Chemiluminescent reactions are sometimes described as generating “cool light”,7 which leads some students to the false impression that these reactions do not generate heat. It is therefore important to demonstrate that chemiluminescent reactions are capable of generating heat. This can be done by estimating12 the enthalpy of the lightstick reaction (eq 1) by using the enthalpies of formation of the compounds given in Table 1 and the generalized equation: Table 1. Enthalpy of Formation of the Compounds Involved in the Lightstick Reaction Figure 2. Energy level diagram depicting possible processes occurring in the lightstick reaction. S0 and S1 (bold horizontal lines) represent the ground and excited electronic states of a dye molecule, respectively. Thin horizontal lines represent vibrational energy levels in each electronic state. Higher-energy vibrational levels of the ground electronic state are extended upward on the left-hand side of S0. Various processes are represented by letters and arrows: chemical excitation (CE, solid up arrows); vibrational relaxation (VR, dashed arrows); internal conversion (IC, wavy arrow); chemiluminescence (CL, solid down arrows).
Compound
ΔrxnH°/(kJ mol−1)
Phenyl oxalate ester, C14H10O4(s) H2O2(l) CO2(g) Phenol, C6H5OH
−540 −188 −394 −165
Δrxn H ° =
° ° − ∑ Δf Hreactants ∑ Δf Hproducts
(4)
A value of −390 kJ mol−1 is calculated, which is substantially exothermic and therefore an increase in temperature should be expected. In fact, an increase in temperature would be expected even in reactions with chemiluminescent quantum yields close to unity. To see why this is so, the energy available from each reacting phenyl oxalate molecule can be compared with the energy released by each photon emitted. Assuming that the lightstick reaction enthalpy (−390 kJ mol−1) is a good estimate for the energy available for light emission, one can calculate that 6.5 × 10−19 J (0.65 aJ) of energy is released for each phenyl
excitation) can be used to promote the dye molecule from its electronic ground state (Figure 2, S0) to an excited electronic state (Figure 2, S1). Notice that the dye molecule can be promoted to any one of several excited vibrational levels in S1. This depends upon the quantity of energy available from the lightstick reaction. Suppose the dye is promoted to the highest vibrational level in S1 (Figure 2, *). First, the dye molecule 911
dx.doi.org/10.1021/ed200328d | J. Chem. Educ. 2012, 89, 910−916
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oxalate molecule that reacts.13 The energy from the molecular reaction will ultimately be used to release a photon or increase the thermal energy of the molecules in the lightstick. Therefore Erxn = Ephoton + Ethermal
photon with a wavelength of 307 nm (Table 2).16 As a result, the lightstick reaction acts as a “chemical black light”, providing enough energy to excite fluorescent dyes that emit color in the visible wavelength range. The different colors of light emitted from lightsticks, thus, depend on the different fluorescent dyes or mixtures of dyes present in a particular lightstick.2 Color-changing lightsticks uniquely demonstrate the interplay between these various aspects of lightsticks. Colorchanging lightsticks contain two dyes: an unstable red-emitting dye and a stable emitter of a different color (Table 3). The
(5)
where Erxn is the energy released by the reaction, Ephoton is the energy of the photon emitted, and Ethermal is the increase in thermal energy of the lightstick system. Even if every reacting phenyl oxalate molecule resulted in an excited dye molecule followed by emission of a photon, there would still be excess energy. This is because the chemical excitation energy (Figure 2, CE) is always greater than or equal to the energy of the photon the dye emits (Figure 2, CL). Also, not all of the energy from this reaction is transferred to the dye; as one example, the two CO2 molecules produced in step 2 of Scheme 1 could carry away reaction energy through translations, rotations, and vibrations. Thus, there will always be some thermal energy. The photonic energy can be determined using the following relation: E=
hc λ
Table 3. Dye Mixtures in Color-Changing Lightsticks Color LightstickType Pink to Blue Orange to Green
Dyes Present Stable blue; Unstable red Stable green; Unstable red
Initial
Intermediate
Final
Pink
Magenta
Blue
Orange
Yellow
Green
unstable dye is susceptible to photobleaching (last step of Scheme 2) as well as chemical oxidation by hydrogen peroxide
(6)
where h is Planck’s constant (6.626 × 10−34 J s), c is the speed of light in a vacuum (3.0 × 108 m s−1), and λ is wavelength of the emitted photon. The thermal energy can then be determined using eq 5. For example, a single blue photon has an energy equivalent of 0.42 aJ (Table 2), leaving 0.23 aJ for
Scheme 2. Possible Reactions of Stable (dye1) and Unstable (dye2) Dyes in Color-Changing Lightsticks
Table 2. Conversion of the Lightstick Reaction Energy to Photon Energy and Thermal Energy of Molecules Photon Emitted
λ/nm
Photon Energy/aJ
Thermal Energy/aJ
Fraction of Available Energy Emitted as Light
None Infrared Red Orange Yellow Green Blue Violet Ultraviolet Lowest λ possible
N/A >700 700 610 575 525 470 400 0.36 0.36 0.32 0.30 0.27 0.23 0.15
