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Why does PCC oxidation stop at the aldehyde?
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Mike Hatlee
Why does PCC oxidation stop at the aldehyde?
The PCC oxidation never gets to the aldehyde. Reaction with an alcohol yields a chromate intermediate. Normally this hydrolyzes to an aldehyde which is immediately oxidized to an acid. The PCC technique is done in an anhydrous medium which prevents hydrolysis. When this step is complete, any excess oxidant is destroyed and, only then is the aldehyde generated.
The PCC oxidation never gets to the aldehyde. Reaction with an alcohol yields a chromate intermediate. Normally this hydrolyzes to an aldehyde which is immediately oxidized to an acid. The PCC technique is done in an anhydrous medium which prevents hydrolysis. When this step is complete, any excess oxidant is destroyed and, only then is the aldehyde generated.
Aluminum has an electronegativity of 1.5. Exactly at the boundary where oxides of elements start becoming acidic.
Oxides of elements of electronegativities below 1.5 readily give off their oxygens in water to grab water’s hydrogens forming hydroxide ions. This is because the bonding is often ionic hence the bond between the element and the oxygen breaks easily in water.
However, beyond 1.5, the element will hold on to the oxygen more strongly and will instead form acids by covalently bonding with more oxygens to form acid anions. The negative charge is going to exist in the oxygen atoms.
Al2O3, being exactly 1.5, is equally likely to give off or grab oxygens because of the de-facto covalent/ionic nature of its bonds. In acidic conditions, protons in the solution will protonate the oxygen forming water and Al’s corresponding salt.
In certain basic coditions however, hydroxide ions promote the formation of more hydroxide ions from Al’s oxygens forming a hydrated aluminate complex and water. The reaction here is way more complicated than a regular acid-base reaction.
This is the simplest, most easily understandable explanation I can give you. I am a mere AS Level student and I don’t know that much about chemistry as much as some other folks around here. Just trying to help out curious people like you in any way I can. ;)
Aluminum has an electronegativity of 1.5. Exactly at the boundary where oxides of elements start becoming acidic.
Oxides of elements of electronegativities below 1.5 readily give off their oxygens in water to grab water’s hydrogens forming hydroxide ions. This is because the bonding is often ionic hence the bond between the element and the oxygen breaks easily in water.
However, beyond 1.5, the element will hold on to the oxygen more strongly and will instead form acids by covalently bonding with more oxygens to form acid anions. The negative charge is going to exist in the oxygen atoms.
Al2O3, being exactly 1.5, is equally likely to give off or grab oxygens because of the de-facto covalent/ionic nature of its bonds. In acidic conditions, protons in the solution will protonate the oxygen forming water and Al’s corresponding salt.
In certain basic coditions however, hydroxide ions promote the formation of more hydroxide ions from Al’s oxygens forming a hydrated aluminate complex and water. The reaction here is way more complicated than a regular acid-base reaction.
This is the simplest, most easily understandable explanation I can give you. I am a mere AS Level student and I don’t know that much about chemistry as much as some other folks around here. Just trying to help out curious people like you in any way I can. ;)
The PCC oxidation never gets to the aldehyde. Reaction with an alcohol yields a chromate intermediate. Normally this hydrolyzes to an aldehyde which is immediately oxidized to an acid. The PCC technique is done in an anhydrous medium which prevents hydrolysis. When this step is complete, any excess oxidant is destroyed and, only then is the aldehyde generated.
The PCC oxidation never gets to the aldehyde. Reaction with an alcohol yields a chromate intermediate. Normally this hydrolyzes to an aldehyde which is immediately oxidized to an acid. The PCC technique is done in an anhydrous medium which prevents hydrolysis. When this step is complete, any excess oxidant is destroyed and, only then is the aldehyde generated.
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Aluminum has an electronegativity of 1.5. Exactly at the boundary where oxides of elements start becoming acidic.
Oxides of elements of electronegativities below 1.5 readily give off their oxygens in water to grab water’s hydrogens forming hydroxide ions. This is because the bonding is often ionic hence the bond between the element and the oxygen breaks easily in water.
However, beyond 1.5, the element will hold on to the oxygen more strongly and will instead form acids by covalently bonding with more oxygens to form acid anions. The negative charge is going to exist in the oxygen atoms.
Al2O3, being exactly 1.5, is equally likely to give off or grab oxygens because of the de-facto covalent/ionic nature of its bonds. In acidic conditions, protons in the solution will protonate the oxygen forming water and Al’s corresponding salt.
In certain basic coditions however, hydroxide ions promote the formation of more hydroxide ions from Al’s oxygens forming a hydrated aluminate complex and water. The reaction here is way more complicated than a regular acid-base reaction.
This is the simplest, most easily understandable explanation I can give you. I am a mere AS Level student and I don’t know that much about chemistry as much as some other folks around here. Just trying to help out curious people like you in any way I can. ;)
Aluminum has an electronegativity of 1.5. Exactly at the boundary where oxides of elements start becoming acidic.
Oxides of elements of electronegativities below 1.5 readily give off their oxygens in water to grab water’s hydrogens forming hydroxide ions. This is because the bonding is often ionic hence the bond between the element and the oxygen breaks easily in water.
However, beyond 1.5, the element will hold on to the oxygen more strongly and will instead form acids by covalently bonding with more oxygens to form acid anions. The negative charge is going to exist in the oxygen atoms.
Al2O3, being exactly 1.5, is equally likely to give off or grab oxygens because of the de-facto covalent/ionic nature of its bonds. In acidic conditions, protons in the solution will protonate the oxygen forming water and Al’s corresponding salt.
In certain basic coditions however, hydroxide ions promote the formation of more hydroxide ions from Al’s oxygens forming a hydrated aluminate complex and water. The reaction here is way more complicated than a regular acid-base reaction.
This is the simplest, most easily understandable explanation I can give you. I am a mere AS Level student and I don’t know that much about chemistry as much as some other folks around here. Just trying to help out curious people like you in any way I can. ;)
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