Potassium Dichromate as an Oxidant

Introduction

Potassium dichromate K₂Cr₂O₇ supplies dichromate ions, Cr₂O₇²⁻, in water. The chromium is formally +6, and in acidic solution dichromate can accept electrons to produce chromium(III). This makes it a useful textbook oxidant, but chromium(VI) compounds have serious health hazards and belong in controlled laboratory handling rather than casual demonstrations.

Core explanation

Assign chromium oxidation state by charge balance. Seven oxygen atoms at −2 total −14. The ion has −2 overall, so two chromium atoms must total +12 and each is +6. Potassium ions in K₂Cr₂O₇ are +1 each and balance the dichromate charge; they are spectators in the essential reduction half-reaction. The formula does not contain free Cr⁶⁺ cations. Chromium(VI) is a formal oxidation-state assignment inside a bonded oxoanion.

In acid, the balanced reduction is Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O. Check atoms: two Cr each side; seven O become seven waters; fourteen H become fourteen H in water. Check charge: left −2 + 14 − 6 = +6, right 2(+3) = +6. Each Cr decreases oxidation state from +6 to +3, gaining three electrons; two Cr atoms account for six electrons. This is the core electron balance behind dichromate oxidations.

To combine with Fe²⁺ oxidation, use Fe²⁺ → Fe³⁺ + e⁻ six times. Electrons cancel, giving Cr₂O₇²⁻ + 14H⁺ + 6Fe²⁺ → 2Cr³⁺ + 7H₂O + 6Fe³⁺. The left charge is −2 + 14 + 12 = +24; the right is +6 + 18 = +24. This equation illustrates that dichromate is reduced while iron(II) is oxidised. Writing “dichromate oxidises iron” refers to its effect on Fe, not to a rise in chromium oxidation number.

Chromate CrO₄²⁻ and dichromate are related by pH-dependent equilibria in solution. In acidic conditions, dichromate or hydrogen chromate can become important; at more alkaline pH, chromate can be favoured. The textbook acidic half-reaction is therefore not a universal equation for every dichromate-containing mixture. Complexation and concentration can also affect speciation and measured potential.

The orange colour often associated with dichromate should not be explained by a simple Cr d–d transition: formal Cr(VI) is d⁰ in the basic count. Charge-transfer absorption involving oxygen and chromium contributes strongly. Cr³⁺ product species can have different colours depending on hydration and ligands. Colour change can support that chemistry occurred, but it is not a complete stoichiometric measurement without analytical controls.

Safety is a substantive difference, not a footnote. Chromium(VI) is toxic and classified as a human carcinogen in occupational exposure guidance. A teaching note should use the ion for equation practice without recommending open handling or disposal down drains. Actual procedures require institutional protocols, suitable protective measures and regulated waste management. Cr(III) and Cr(VI) should not be treated as toxicologically interchangeable simply because both contain chromium.

Step-by-step reasoning

1. Calculate Cr oxidation state in Cr₂O₇²⁻. 2. Write two Cr³⁺ products for the two Cr atoms. 3. Balance O with seven H₂O and H with fourteen H⁺. 4. Add six electrons on the reactant side and verify charge. 5. Multiply a partner oxidation half-reaction to cancel electrons.

Visual explanation

Draw a dichromate ion with two Cr labels “+6,” then an arrow to two Cr³⁺ labels. Under the arrow write “six electrons accepted total.” Place oxygen and hydrogen balances below it as seven waters and fourteen protons.

Real-world analogy

Two workers each accept three tokens; the pair receives six tokens altogether. Each Cr(VI) gains three electrons to become Cr(III), so the dichromate unit accepts six electrons in the acidic half-reaction.

Real-world example

Classical redox titrations used dichromate as a standard oxidant for reducing agents such as Fe²⁺ under carefully controlled acidic conditions. Modern safety practice treats Cr(VI) exposure and waste as significant concerns.

Why?

Why are six electrons required in the reduction half-reaction? Two chromium atoms each move from formal +6 to +3, a decrease of three oxidation-number units per atom, for six electrons overall.

Common misconception

“K₂Cr₂O₇ contains free Cr⁶⁺ ions.” It contains K⁺ and the bonded dichromate oxoanion. Chromium(VI) is a formal state within Cr₂O₇²⁻, not a bare aqueous cation.

Worked example

Balance dichromate oxidising Fe²⁺ in acid. Start with Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O. Multiply Fe²⁺ → Fe³⁺ + e⁻ by six and add. The net ionic equation is Cr₂O₇²⁻ + 14H⁺ + 6Fe²⁺ → 2Cr³⁺ + 7H₂O + 6Fe³⁺. Both atom and charge balances match.

Quick check

1. What is the formal chromium oxidation state in Cr₂O₇²⁻? Answer: +6 for each chromium atom.

Exam focus

Use the acidic half-reaction only when the medium supports it. Show electron and charge checks, and distinguish formal Cr(VI) in oxoanion from free cation language. Treat Cr(VI) hazards as materially different from a routine salt.

Advanced insight

OpenStax gives the acidic dichromate half-reaction in its transition-metal chapter, https://openstax.org/books/chemistry-2e/pages/19-1-occurrence-preparation-and-properties-of-transition-metals-and-their-compounds. OSHA summarises Cr(VI) occupational hazards at https://www.osha.gov/chromium.

Summary

Dichromate contains two formal Cr(VI) atoms and accepts six electrons to form two Cr³⁺ ions in the standard acidic half-reaction. Medium alters speciation, and Cr(VI) toxicity requires controlled handling and disposal.

Practice questions

1. How many electrons does one dichromate ion accept when both Cr(VI) become Cr(III)? Answer: Six. 2. How many H⁺ appear in the acidic half-reaction? Answer: Fourteen. 3. Is dichromate oxidised or reduced when it converts Fe²⁺ to Fe³⁺? Answer: Reduced; it acts as the oxidant. 4. Does the orange colour prove a d–d transition from Cr(VI)? Answer: No. Formal Cr(VI) is d⁰, and charge-transfer absorption is important.