All metals form basic oxides. Metallic oxides are basic in character.
Polonium is a metal, so it should form basic oxides but it forms acidic oxides. Why are its oxides acidic in character?
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3$\begingroup$ Polonium is actually a metalloid, and not strictly speaking classified as a metal… $\endgroup$– F'xCommented Oct 5, 2012 at 7:34
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$\begingroup$ Mn2O7 and CrO3 are acidic oxides of metals. $\endgroup$– PoutnikCommented Jul 15 at 3:48
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$\begingroup$ The oxide of polonium in its preferred oxidation state ($\ce{PoO2}$; +4) is predominantly basic, but amphoteric if dissolved in concentrated aqueous alkali, or fused with potassium hydroxide in air. $\endgroup$– Nilay GhoshCommented Jul 15 at 4:05
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4 Answers
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All metals form basic oxides.
That's true. I cannot remember a metal that does not form basic oxides.
Metallic oxides are basic in character.
Not necessarily. Polonium is not the only metal that forms acidic oxides, Mn, Cr and some others do so as well.
An oxide is considered acidic when it forms a hydroxide that is acidic (well, this is actually a simplification, but let's stick with it for a while). A hydroxide is acidic when it can lose a proton, giving a stable anion. The acidity of the hydroxide and metallicity of the element formed is not connected in any way. The acidity depends on the ability of the remaining parts of the molecule to stabilize negative charge, which usually correlates with the $\ce{O/OH}$ ratio and the electronegativity of the central atom. So, $\ce{H6TeO6}$ is quite a weak acid and $\ce{H2CrO4}$ or $\ce{HMnO4}$ are quite strong acids.
answered Sep 24, 2012 at 11:06
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$\begingroup$ Isn't O^2- extremely unstable? I mean I can understand why O would want 2 more electrons but it doesn't really want to be ionic and for example Mg^2+ O^2- would be an ionic oxide(an oxide since O is more electronegative than Mg and ionic since the 2 have opposite charges) and oxygen wants to form covalent bonds like it does in Iron Oxide, not ionic bonds like a chloride would so ionic bonds to oxygen anions or cations are extremely unstable $\endgroup$– CatersCommented Jul 28, 2014 at 20:47
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2$\begingroup$ @caters It isn't related to the topic, but in general there is no purely ionic or covalent compounds. Some oxides, like $\ce{MgO}$, have significant ionic bonding, but just as well there is no chloride that does not have significant covalent bonding. In addition, there is so known lattice energy, an energy, freed as result of bonding of ions into lattice. It is quite big and may be enough to stabilize some unstable ions. $\endgroup$ Commented Jul 29, 2014 at 5:50
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$\begingroup$ MgCl2 has ionic bonding, not covalent. Same with NaCl. These are all ionic chlorides. $\endgroup$– CatersCommented Jul 31, 2014 at 11:47
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2$\begingroup$ @caters CsCl, most ionic of all chlorides, has only 75% of ionic bonding and 25% of covalent bonding. $\endgroup$ Commented Jul 31, 2014 at 12:44
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$\begingroup$ MgCl2 has significant covalent character, as can be seen from it's layer structure (acc. to JD Lee)... no time to elaborate on this, you can search yourself maybe $\endgroup$ Commented Dec 1, 2019 at 16:05
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Another way of looking at this is to understand that the structure is always
$$\ce{M-O-H}$$
where $\ce{M}$ is the metal atom. If the $\ce{O-H}$ bond is weaker than the $\ce{M-O}$ bond, the material acts as an acid. Otherwise it acts as a base.
What determines the relative strength is a bit complex. Suffice it to say that it depends on oxidation state of the $\ce{M}$ atom as well as the constituents in the rest of the molecule and the environment all play a role.
Some materials are amphoteric. That is they can split either way. A typical example is aluminum hydroxide, $\ce{Al(OH)_3}$, which has three $\ce{O-H}$ bonds. It is usually a base, but under the right conditions it breaks the other way and is a weak acid called aluminic acid. The formula for that is usually written as $\ce{H3AlO_3}$, but the structure is the same.
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1$\begingroup$ not all metal oxides are this way. For example iron oxides have this structure: FeOx where x is the number of oxygen atoms. No hydrogens here. and yes oxides do tend to be basic When there are hydrogens in the oxide. Otherwise they tend to be acidic(which is how come CO2 acidifies water instead of making it more basic) $\endgroup$– CatersCommented Jul 28, 2014 at 20:51
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While virtually all metals form a basic oxide or amphoteric (reacts with both bases and acids) oxides like alumina and both Tin(II) and Tin(IV), some also form acidic oxides like Chromium(VI) oxide. It is better to describe oxides according to their reactivity with acids and bases as you are not trying to convert them to the hydrates/hydroxides.
It isn't incredible for Polonium to form only an acidic oxide.
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$\begingroup$ It does form basi oxide: en.wikipedia.org/wiki/Polonium_monoxide $\endgroup$– MithoronCommented Apr 22, 2015 at 13:41
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$\begingroup$ @Mithoron I do not read this from Wikipedia. It seems to say only that Po(II) oxide or hydroxide is oxidized to Po(IV) (presumably displacing hydrogen from the water). Only Po(IV) oxide seems to be well characterized, and that compound's WP article actually calls it amphoteric. $\endgroup$ Commented Jul 16 at 1:26
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Strictly speaking, "polonium oxide" is not necessarily acidic. There are three known polonium oxides, but only the most stable one, $\ce{PoO2}$, is well charactetized. Wikipedia reports that this oxide is actually amphoteric:
When polonium dioxide is hydrated, polonous acid ($\ce{H2PoO3}$), a pale yellow, voluminous precipitate, is formed. Despite its name, polonous acid is an amphoteric compound, reacting with both acids and bases.[1][2]
That polonium dioxide is amphotetic is not really remarkable because oxides with the other element in the +4 oxidation state run the gamut from acidic (sulfur dioxide) to basic (thorium dioxide) with some oxides like titanium dioxide showing amphoteric character in between. Incidentally, silicon dioxide is one of the few nonmetal oxides that qualify as showing a limited amphoteric character, reacting as a proton acceptor with hydrofluoric acid.
Cited References
Holleman, Arnold Frederik; Wiberg, Egon (2001), Wiberg, Nils (ed.), Inorganic Chemistry, translated by Eagleson, Mary; Brewer, William, San Diego/Berlin: Academic Press/De Gruyter, p. 594, ISBN 0-12-352651-5 Bagnall, K. W.; D'Eye, R. W. M. (1954).
Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. p. 780. ISBN 978-0-08-037941-8.
