Weird Science

Hungry Liquid

Polish ver­sion is here

The fol­lo­wing article was ori­gi­nally publi­shed in the jour­nal for edu­ca­tors Che­mia w Szkole (eng. Che­mi­stry in School) (4/2024):

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Ples M., Żar­łoczny roz­twór (eng. Hun­gry Liquid), Che­mia w Szkole (eng. Che­mi­stry in School), 4 (2024), Agen­cja AS Józef Szew­czyk, pp. 36–37

A See­min­gly Impos­si­ble Para­dox

Can an object be pla­ced inside a con­ta­i­ner if its size exce­eds the con­ta­i­ner’s capa­city? It seems impos­si­ble.

Or is it? Let’s take a clo­ser look. After some con­si­de­ra­tion, we might rea­lize that in order to accom­plish this feat, we sim­ply need to rear­range the par­tic­les of the "hid­den" object. The task, altho­ugh the­o­re­ti­cally pos­si­ble, still seems extre­mely dif­fi­cult. This is where che­mi­stry comes to our aid.

Thanks to elec­tro­che­mi­stry, we can easily pack a rela­ti­vely large piece of metal into a small bea­ker.

What Do We Need?

For­tu­na­tely, the requ­i­red sub­stan­ces are easy to obtain:

Cop­per(II) sul­fate, as a heavy metal com­po­und, can be hazar­dous. Stan­dard labo­ra­tory pre­cau­tions sho­uld be fol­lo­wed when han­dling che­mi­cals.

We have alre­ady enco­un­te­red hydra­ted cop­per(II) sul­fate in a humo­rous but suc­cess­ful attempt to "weigh" its color [1]. This salt usu­ally appe­ars as well-for­med, bri­ght blue cry­stals (Fig.1).

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Fig.1 – Cry­stals of hydra­ted cop­per(II) sul­fate

Hou­se­hold alu­mi­num foil works per­fec­tly as a source of alu­mi­num. Obta­i­ning sodium chlo­ride is even sim­pler, as it is a com­mon kit­chen ingre­dient.

Expe­ri­ment

We need to pre­pare 100 cm³ (3.4 fl oz) of a 10% cop­per(II) sul­fate solu­tion. The liquid has a beau­ti­ful blue color (Fig.2).

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Fig.2 – Cop­per(II) sul­fate solu­tion

Roll a piece of alu­mi­num foil into a tube appro­xi­ma­tely 30 cm (12 inches) long. Its dia­me­ter sho­uld be sli­gh­tly smal­ler than that of the bea­ker (Fig.3).

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Fig.3 – Pro­perly pre­pa­red alu­mi­num foil

As seen in Fig.4, the rol­led alu­mi­num foil is indeed much lar­ger than the bea­ker.

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Fig.4 – Expe­ri­men­tal setup

Sub­mer­ging the lower part of the alu­mi­num roll into the cop­per sul­fate solu­tion does not cause any visi­ble chan­ges. Howe­ver, after adding a few cm³ (0.1 fl oz) of a satu­ra­ted room-tem­pe­ra­ture sodium chlo­ride solu­tion, the reac­tion begins almost imme­dia­tely. The alu­mi­num dis­so­lves vigo­ro­u­sly, rele­a­sing heat and incre­a­sing the liquid’s tem­pe­ra­ture. After a few minu­tes, nearly all of the silver-colo­red alu­mi­num has dis­so­lved, while a cer­tain amo­unt of red­dish-brown pre­ci­pi­tate appe­ars in the bea­ker (Fig.5).

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Fig.5 – Final stage of the expe­ri­ment; alu­mi­num nearly com­ple­tely dis­so­lved

Since only a small amo­unt of water vapor has esca­ped the bea­ker, we have achie­ved our goa­l—we suc­cess­fully fit a large piece of alu­mi­num into a small bea­ker!

Expla­na­tion

To under­stand this pro­cess, we need to con­si­der the elec­tro­che­mi­cal series of metals.

The elec­tro­che­mi­cal series of metals, also known as the reac­ti­vity series, ranks che­mi­cal ele­ments with metal­lic pro­per­ties based on their stan­dard elec­trode poten­tial E0 [2] [3]. The refe­rence point in this ran­king is the hydro­gen elec­trode, whose stan­dard poten­tial is arbi­tra­rily set at zero. A sim­pli­fied ver­sion of the elec­tro­che­mi­cal series, arran­ged in incre­a­sing order of E0, is shown below:

Li < K < Na < Ca < Mg < Al < Mn < Zn < Cr < Fe < Cd < Co < Ni < Sn < Pb < H < Sb < Bi < Cu < Ag < Hg < Pt < Au

Hydro­gen, which does not exhi­bit metal­lic pro­per­ties under stan­dard con­di­tions, is high­li­gh­ted in red. It is impor­tant to remem­ber that the lower the stan­dard poten­tial, the more che­mi­cally active the metal. Indeed, the most reac­tive metal­s—such as lithium (Li) and potas­sium (K)—are gro­u­ped on the left, while the noble metals, inc­lu­ding silver (Ag), pla­ti­num (Pt), and gold (Au), are on the right. Ano­ther key prin­ci­ple is that, in most cases, a more reac­tive metal displa­ces a less reac­tive one from solu­tion, while itself ente­ring the solu­tion as cations.

Alu­mi­num is signi­fi­can­tly more reac­tive than cop­per, so it sho­uld rea­dily displace cop­per ions from the solu­tion. Why, then, do we not obse­rve any reac­tion at first? The answer lies in pas­si­va­tio­n—a­lu­mi­num is so reac­tive that even at room tem­pe­ra­ture, it forms a thin, invi­si­ble, imper­me­a­ble layer of oxi­des and hydro­xi­des on its sur­face. This pro­tec­tive layer pre­vents cop­per(II) ions from rea­ching the metal­lic alu­mi­num.

The situ­a­tion chan­ges dra­ma­ti­cally when chlo­ride ions from the dis­so­cia­tion of sodium chlo­ride are intro­du­ced. These chlo­ride ions break down the pro­tec­tive oxide layer, allo­wing the reac­tion to begin: alu­mi­num dis­so­lves, for­ming inso­lu­ble com­po­unds, while cop­per ions are redu­ced to metal­lic cop­per [4].

Al + Cu2+ → Al3+ + Cu↓

It is worth noting that alu­mi­num is not only oxi­di­zed but also under­goes hydro­ly­sis in the solu­tion, lea­ding to the for­ma­tion of com­plex hydro­xide spe­cies. Addi­tio­nally, the cop­per depo­sits as a fine-gra­i­ned pre­ci­pi­tate, gra­du­ally for­ming visi­ble clu­sters at the bot­tom of the bea­ker.

Refe­ren­ces

All pho­to­gra­phs and illu­stra­tions were cre­a­ted by the author.

The above text inc­lu­des minor edi­to­rial modi­fi­ca­tions com­pa­red to the ver­sion publi­shed in the jour­nal, aimed at sup­ple­men­ting and adap­ting it for online pre­sen­ta­tion.

Marek Ples

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