Science
Self Heating Food and Drink Cans
Heater meals – originally developed for military use – are ready-made self-heating meal packs. They can be heated in many ways – by pressing a button on the packaging, unwrapping and shaking the pack, or pouring the contents of one bag into another and waiting for a few minutes – all of which use exothermic chemical reactions. These meals can be used to motivate students to study such reactions relatively safely and without the use of a burner. Plus there is the added value of discussing the negative ecological aspects of disposable meals. For the following experiment, we use the Crosse & Blackwell heater-meal system, which relies on the reaction of magnesium and salt water to produce hydrogen:
Mg (s) + 2H2O (l) -> Mg2+ (aq) + H2 (g) ↑ + 2OH- (aq) s: solid; l: liquid; g: gaseous; aq: in solution; the vertical arrow indicates that gas is released.
This reaction is very slow, due to passivation, so to speed it up, iron and salt are added. Passivation is the process by which a material is made less reactive, usually by the deposition of a layer of oxide on its surface: if you place a strip of magnesium into cold water, its surface will oxidise to magnesium hydroxide (Mg(OH)2), and this coating will prevent further reaction.
Therefore, in the heater meal, iron is added to the magnesium, leading to the production of a local cell – small-scale corrosion that happens where two metals of different reactivity are in contact under humid conditions – which speeds up the exothermic reaction. Because the electron potential of magnesium is lower than that of iron (the less reactive metal), electrons will pass from the magnesium to the iron, and only from there into the water. Although magnesium cations (Mg2+) and hydroxide anions (OH-) continue to be formed, they are separated by the iron and cannot combine to form magnesium hydroxide. As a result, the magnesium does not become passivated by a coating of magnesium hydroxide, which
Heater meals – originally developed for military use – are ready-made self-heating meal packs. They can be heated in many ways – by pressing a button on the packaging, unwrapping and shaking the pack, or pouring the contents of one bag into another and waiting for a few minutes – all of which use exothermic chemical reactions. These meals can be used to motivate students to study such reactions relatively safely and without the use of a burner. Plus there is the added value of discussing the negative ecological aspects of disposable meals. For the following experiment, we use the Crosse & Blackwell heater-meal system, which relies on the reaction of magnesium and salt water to produce hydrogen:
Mg (s) + 2H2O (l) -> Mg2+ (aq) + H2 (g) ↑ + 2OH- (aq) s: solid; l: liquid; g: gaseous; aq: in solution; the vertical arrow indicates that gas is released.
This reaction is very slow, due to passivation, so to speed it up, iron and salt are added. Passivation is the process by which a material is made less reactive, usually by the deposition of a layer of oxide on its surface: if you place a strip of magnesium into cold water, its surface will oxidise to magnesium hydroxide (Mg(OH)2), and this coating will prevent further reaction.
Therefore, in the heater meal, iron is added to the magnesium, leading to the production of a local cell – small-scale corrosion that happens where two metals of different reactivity are in contact under humid conditions – which speeds up the exothermic reaction. Because the electron potential of magnesium is lower than that of iron (the less reactive metal), electrons will pass from the magnesium to the iron, and only from there into the water. Although magnesium cations (Mg2+) and hydroxide anions (OH-) continue to be formed, they are separated by the iron and cannot combine to form magnesium hydroxide. As a result, the magnesium does not become passivated by a coating of magnesium hydroxide, which