Group III Cations: Hydroxides in Buffered Ammonia

Introduction

In the chosen classical scheme, iron(III), aluminium and chromium(III) can be separated as hydroxides under a controlled ammonia/ammonium medium. The purpose of the buffer is to establish a useful OH⁻ level, not to make the solution as basic as possible. Their hydroxides have low solubilities under those conditions, while other ions may remain dissolved for later groups. Colour helps, but mixed solids and variable oxidation state require caution.

Core explanation

The formal ionic equations are Fe³⁺ + 3OH⁻ → Fe(OH)₃(s), Al³⁺ + 3OH⁻ → Al(OH)₃(s), and Cr³⁺ + 3OH⁻ → Cr(OH)₃(s). Each trivalent cation needs three hydroxides for charge balance. Their Ksp expressions are [M³⁺][OH⁻]³ in a simplified concentration model. Because OH⁻ is cubed, relatively small pH changes can have a large effect on the ion product.

Ammonia is a weak base: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻. Adding NH₄⁺ as part of a buffered medium suppresses uncontrolled NH₃ ionization through a common-ion effect and helps hold pH within a useful range. The NH₃/NH₄⁺ ratio determines approximate pH, while total buffer amount influences resistance to disturbances. The resulting free OH⁻ can precipitate low-Ksp trivalent hydroxides without automatically precipitating every divalent hydroxide under the scheme's intended concentrations.

Fe(OH)₃ is often described as rust-brown or reddish brown, Al(OH)₃ as white, and Cr(OH)₃ as grey-green or greenish. These descriptions are not perfect fingerprints: hydrated forms, particle size and mixed precipitates can change appearance. Iron(II), if present, has different chemistry and may oxidize in air, creating brown Fe(III)-containing material that confuses a simple colour reading. A cation assignment should use a confirmatory response beyond shade.

Aluminium and chromium hydroxides can be amphoteric, dissolving under suitable excess strong base through hydroxo species. However, an ammonia buffer is not the same as excess strong sodium hydroxide. The terms “base” and “excess base” must specify actual pH and ligand chemistry. A hydroxide may precipitate in buffered ammonia yet dissolve under a different strong-base condition; this is not inconsistent.

Metal hydrolysis complicates a bare M³⁺ picture. Highly charged hydrated metal ions can donate protons from coordinated water before a visible bulk hydroxide forms. The textbook solid formula M(OH)₃ is a useful composition shorthand, while real precipitates may be hydrated or poorly crystalline. This is another reason observations need chemical confirmation and controlled conditions.

As always, separate the Group III solid from the filtrate before later sulfide or carbonate reasoning. A green solid might contain chromium hydroxide, but it could also include another species retained from an earlier stage if separation was incomplete. Phase history matters as much as colour.

The group can be understood from onset thresholds: [OH⁻]onset = (Ksp/[M³⁺])^(1/3). Different metal concentrations alter the pH at which each solid appears. Grouping is therefore a practical window, not an immutable law that all three appear at exactly the same pH.

Step-by-step reasoning

1. Identify the filtrate entering the buffered-hydroxide stage. 2. Relate NH₃/NH₄⁺ to a controlled OH⁻ level. 3. Write M(OH)₃ formulas and Ksp expressions for trivalent ions. 4. Record precipitate colours as clues only. 5. Separate and use further chemistry to confirm Fe³⁺, Al³⁺ or Cr³⁺.

Visual explanation

Draw a pH gauge held in a middle region by NH₃/NH₄⁺. Mark Fe(OH)₃, Al(OH)₃ and Cr(OH)₃ crossing their Ksp thresholds while later-group ions remain on the dissolved side.

Real-world analogy

A thermostat holds temperature in a narrow range so selected materials change state while others do not. The ammonia buffer similarly holds an OH⁻ window that makes a selective precipitation possible.

Real-world example

A reddish-brown hydroxide fraction after the buffered stage is consistent with Fe³⁺, while a white component may suggest Al³⁺. A mixture can show a combined colour and needs a separate confirmation for both.

Why?

Why use NH₃ together with NH₄⁺ instead of assuming unlimited OH⁻? The conjugate pair moderates pH and free hydroxide, preserving a separation window rather than indiscriminately precipitating additional ions.

Common misconception

“All coloured Group III precipitate is pure Fe(OH)₃.” Chromium can contribute greenish material, and mixtures or oxidation can alter colour; independent confirmation is needed.

Worked example

A hypothetical trivalent metal ion has Ksp(M(OH)₃) = 10⁻³³ and free [M³⁺] = 10⁻³ M. Its onset OH⁻ is (10⁻³³/10⁻³)^(1/3) = 10⁻¹⁰ M. At 25 °C this corresponds to pOH 10 and pH about 4. The arithmetic illustrates the cube-root threshold; actual Fe, Al or Cr speciation may include hydrolysis and complexes, so their observed onset needs the full model.

Quick check

1. Why is hydroxide concentration cubed in Fe(OH)₃'s Ksp expression? Answer: One formula unit releases three OH⁻ ions, giving the factor [OH⁻]³ by dissolution stoichiometry.

Exam focus

Explain the NH₃/NH₄⁺ control and use correct trivalent formulas. Treat colour and group number as clues, not proof.

Advanced insight

Ammonia may also coordinate some metal ions, changing free metal concentration. A precise separation model can require both hydroxide and ammine-complex equilibria in addition to the buffer equations.

Summary

Fe³⁺, Al³⁺ and Cr³⁺ can precipitate as low-solubility hydroxides in a controlled ammonia/ammonium medium. The buffer sets free OH⁻, and Ksp sets onset. Colour, hydrolysis, amphoterism and complexation limit simple one-test identification.

Practice questions

1. What is the formal hydroxide formula for Cr³⁺? Answer: Cr(OH)₃. 2. Which pair controls pH in buffered ammonia? Answer: NH₃ and NH₄⁺. 3. Why is a white Group III solid not uniquely Al(OH)₃? Answer: Other white solids or mixed fractions may be present, and appearance alone is not conclusive. 4. What changes if free M³⁺ concentration increases at fixed Ksp? Answer: The OH⁻ onset threshold decreases, so precipitation can begin at a lower hydroxide level.