Amphoteric metals with alkalis

Every metal you have studied in this chapter reacts with a dilute acid to give hydrogen — that is the general rule. What is unusual is that three of those metals also react with a strong alkali and give hydrogen there too. This double behaviour, reacting with both an acid and a strong alkali to release the same gas, is confined to aluminium, zinc, and lead among the Class 9 syllabus metals.

The three defining reactions. Each metal displaces hydrogen from sodium hydroxide (or potassium hydroxide) solution, and a soluble salt of the metal forms alongside the gas:

where: $NaAlO_2$ is sodium aluminate (a soluble salt), and water is a co-reactant needed to balance aluminium's reaction — this is the one equation of the three that consumes water on the reactant side.

where: $Na_2ZnO_2$ is sodium zincate, formed directly from zinc and NaOH with no extra water needed.

where: $Na_2PbO_2$ is sodium plumbite. Note the "-ite" ending here rather than "-ate" — lead's salt is conventionally named plumbite, not plumbate, at this level. Swap NaOH for KOH in any of the three equations and you get the matching potassium salt — potassium aluminate, potassium zincate, potassium plumbite — with hydrogen liberated exactly as before.

Who is in the group and who is not. Only aluminium, zinc, and lead behave this way among the metals of this chapter. Magnesium, iron, sodium, potassium, and calcium do not liberate hydrogen from NaOH or KOH — they may react with acids perfectly normally, but they stay unreactive toward alkalis under these conditions. So "reacts with acid" and "reacts with alkali" are two separate properties of a metal, and only three metals in this syllabus happen to have both.

Why examiners bother naming this at all. It connects two ideas you already know — the activity series (which ranks how readily metals lose electrons) and the acid-hydrogen test — into a slightly wider picture: a metal's position in the activity series does not by itself predict whether it will also react with an alkali. Al, Zn, and Pb sit in the middle of the series, reactive enough to attack alkalis under these conditions, while stronger reducing metals like sodium and potassium, and weaker ones like copper, do not show this particular alkali reaction.

Understanding the reaction is straightforward at this level: NaOH and KOH are strong alkalis, and Al, Zn, and Pb are metals whose oxide layer or surface can be attacked by such a strong alkali just as it is attacked by a strong acid, so the metal dissolves into solution while hydrogen is displaced from the water/hydroxide system. Step 1. In each case the metal loses electrons to the surrounding solution, exactly as it does with dilute acid — this is still an ordinary reaction where the metal is oxidised. Step 2. The hydrogen ions needed to form $H_2$ gas come from the water/hydroxide medium rather than from a free acid, which is why the alkali equations show water or extra $OH^-$ appearing in the reactants rather than a simple acid. Step 3. The metal ion produced does not stay as a simple salt like a chloride or sulphate — it combines with the alkali metal and oxygen to give a distinct "-ate"/"-ite" compound (aluminate, zincate, plumbite), which is what makes these products look different from the familiar salts formed with acids. Step 4. Only Al, Zn, and Pb are reactive enough, and their compounds stable enough, for this particular alkali attack to happen at the Class 9 level — the complete chemical reason (the formal idea of an "amphoteric" oxide or hydroxide, one that reacts with both acids and bases) is reserved for Class 10 Chapter 3, so for now you memorise the pattern rather than derive it from first principles.

Question. A student is given three unlabelled metal strips known to be magnesium, aluminium, and lead, plus a bottle of sodium hydroxide solution and a bottle of dilute hydrochloric acid. Using only these alkali and acid tests, how can the student identify which strip is magnesium?

Step 1. Test each strip with dilute HCl first. All three metals — magnesium, aluminium, and lead — react with a dilute acid to liberate hydrogen with a "pop" on a lit splint, so this test alone cannot distinguish magnesium from the other two.

Step 2. Test each strip with sodium hydroxide solution instead. Recall the special group: only aluminium, zinc, and lead liberate hydrogen from NaOH; magnesium does not.

Step 3. Observe the three strips in NaOH. Two of them show bubbling (effervescence) and give a confirmatory pop test, following $2Al + 2NaOH + 2H_2O \rightarrow 2NaAlO_2 + 3H_2\uparrow$ for the aluminium strip and $Pb + 2NaOH \rightarrow Na_2PbO_2 + H_2\uparrow$ for the lead strip. One strip shows no visible reaction at all in the NaOH.

Step 4. Identify the strip. The strip that reacted with the acid (Step 1) but showed no reaction with the alkali (Step 3) must be magnesium, since magnesium reacts with acids but is not among the alkali-reacting metals (Al, Zn, Pb).

• Assuming that reacting with acid automatically means a metal also reacts with alkali — only aluminium, zinc, and lead show both; magnesium, iron, sodium, potassium, and calcium react with acids but not with NaOH/KOH in this way.
• Misnaming the salt products as ordinary chlorides or sulphates — the alkali reaction always gives an "-ate"/"-ite" salt: sodium aluminate, sodium zincate, sodium plumbite (note "-ite", not "-ate", for lead).
• Forgetting the extra water molecules needed to balance aluminium's equation — $2Al + 2NaOH + 2H_2O \rightarrow 2NaAlO_2 + 3H_2\uparrow$ is the one of the three that needs water written on the reactant side; zinc and lead do not.
• Reaching for the formal "amphoteric oxide/hydroxide" explanation — that theoretical treatment belongs to Class 10; at this level, simply recognise and correctly write the reaction for these three named metals.

Ask the tutor: Want to see how these same three metals' alkali reactions connect to the fuller "amphoteric oxide" theory you will meet in Class 10?