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Hydrated metals form acidic solutions but metal hydroxides are basic Watch

    • Thread Starter

    I'm learning online and struggling with this. Very grateful for any help given.

    I understand that adding a Group 1 or Group 2 metal to water usually produces a metal hydroxide solution (to varying degrees) and hydrogen. This metal hydroxide solution will be basic. As I understand it, the metal ions act as reducing agents by causing water molecules to become hydrogen gas and hydroxide ions. The increased concentration of hydroxide ions causes the pH of the solution to rise - hence the term alkali/alkali metals.

    I also understand that adding some metal ions to water causes a solution that is acidic. I understand that this is because the hydrated ion acts as a Lewis acid by accepting lone pairs from oxygen atoms on water molecules, forming a hydrated metal ion complex. Depending on the charge density of the metal ion, the outer hydrogen atoms are left electron-deficient and exposed. They are plucked off by water molecules to form oxonium ions, causing the pH of the solution to fall. This also causes the hydrated metal ion complex to become a metal hydroxide precipitate. I believe this may or may not become a neutral precipitate.

    What I don't understand is this:
    - How can a metal hydroxide solution be basic in the first set of circumstances but acidic in the second?
    - Is the second solution not basic because the hydroxide ions are not actually in solution but bonded to the metal cation?
    - Why, when a metal is added to water and an ionic solution containing hydroxide ions and metal cations is formed, do hydrated metal ion complexes not form, causing an acidic solution? Is it because there are not enough water molecules to surround the metal cations (because they have all been turned into hydroxide ions)? Or is it because the hydroxide ions are a stronger base and don't allow the water molecules be in close proximity to water molecules?

    Again, thank you for taking the time to read and help. This has been a long-term problem and I would love to know the answer.

    When you put the metal into water, it reacts with the water, e.g. Ba + 2H2O -> Ba(OH)2 + H2, the Ba(OH)2 then dissociates releasing OH-, i.e. it is alkaline.

    If you react the metal with something first, then dissolve that compound with water, the H in water isn't going to get reduced and OH- won't be made.

    Instead the metal ion, e.g. Al3+ will be hydrated with 6H2O. Since the Al has such high charge density, it pulls on the e- in the water molecules, weakening the O-H bonds, allowing a H+ to be pulled off by a nearby water molecule, i.e. it is acidic.

    The key to it is recognising when the oxidation of the metal took place.
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