Chemistry
How Chromium(VI) Compounds Act as Strong Oxidants in Acidic Conditions
Quick fact
In acidic solution, the standard reduction potential of chromium(VI) to chromium(III) is about +1.33 V, which is even stronger than chlorine gas (E° = +1.36 V). This means acidic dichromate can oxidize many organic compounds, including ethanol to acetaldehyde.
Why this is interesting
You may have seen orange dichromate solution turn green when mixed with an alcohol. Why does this happen, and why does it only work when the solution is acidic?
Read the full explanation
Understanding How Chromium(VI) Compounds Act as Strong Oxidants in Acidic Conditions
Chromium(VI) compounds are strong oxidants because chromium in its +6 oxidation state has a high tendency to gain electrons. In acidic conditions, this tendency is amplified. Imagine the Cr(VI) species as a hungry electron acceptor: it wants to return to a more stable, lower oxidation state (Cr(III)). In basic conditions, the oxidizing power is much weaker. The key is that in acidic medium, the oxygen atoms on the chromium species can combine with hydrogen ions (protons) to form water, which pulls the reaction forward. For example, in the dichromate ion (Cr₂O₇²⁻), when it is reduced, it gains electrons (e⁻) and hydrogen ions (H⁺) to become two chromium(III) ions (Cr³⁺) and water. The presence of acid provides those protons, making the reaction favorable. Without acid, the oxidizing strength drops significantly. This is why you might see dichromate used as a test for alcohols: when you add an alcohol to an acidic dichromate solution, the alcohol is oxidized (loses electrons), while the dichromate is reduced, turning from orange (Cr(VI)) to green (Cr(III)).
A deeper explanation
The root cause of the strong oxidizing power is a combination of electron affinity and energetics. Chromium(VI) has a high positive oxidation state, meaning it has 'lost' many electrons and is eager to regain them. In acidic solution, the reduction half-reaction for dichromate is: Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O (E° = +1.33 V) The Nernst equation shows that the actual potential (E) increases with increasing concentration of H⁺. Specifically, E = E° - (0.0592/n) log(Q), where Q = [Cr³⁺]²/([Cr₂O₇²⁻][H⁺]¹⁴). Because H⁺ concentration is in the denominator of the reaction quotient, a higher [H⁺] (more acidic) makes Q smaller, making the log term more negative, so E becomes more positive. Therefore, the oxidizing power is enhanced. Additionally, protons are consumed in the reaction, tying the redox process to acid-base chemistry. The electron transfer changes the coordination of chromium: in Cr(VI) it is tetrahedral (as in dichromate), while in Cr(III) it is octahedral (as the hydrated ion [Cr(H₂O)₆]³⁺). The driving force includes the high stability of the Cr(III) species and the formation of water. In basic conditions, the corresponding half-reaction is: CrO₄²⁻ + 4H₂O + 3e⁻ → Cr(OH)₃ + 5OH⁻ (E° = -0.13 V) which is far less positive, explaining why Cr(VI) is a much weaker oxidant in alkaline media. This pH dependence is exploited in organic chemistry: acidic dichromate or chromic acid is used for the controlled oxidation of primary alcohols to aldehydes (like the Jones oxidation) and secondary alcohols to ketones. The reaction with a primary alcohol can be stopped at the aldehyde stage, but if water is present, further oxidation to a carboxylic acid occurs. This illustrates how the strong oxidant can perform multiple oxidations. Understanding this concept also helps explain the environmental persistence of Cr(VI) in acidic groundwater: it remains capable of oxidizing organic matter and is highly toxic, whereas reduction to Cr(III) is less harmful.