Did you know that the same metal—iron—can form aqueous ions with very different acid strengths just by changing its oxidation state? A hydrated iron(III) ion, [Fe(H₂O)₆]³⁺, is substantially more acidic than a hydrated iron(II) ion, [Fe(H₂O)₆]²⁺. The iron(III) centre has a higher positive charge and greater polarising power. That makes water molecules coordinated to Fe³⁺ more willing to transfer a proton to surrounding water, generating H₃O⁺.
For JC1 and JC2 students searching for an H2 Chemistry tutor in Hougang, Fe3+ vs Fe2+ aqua ion acidity, iron(III) hydrolysis, charge density and pH, hexaaquairon complex ions, transition metal hydroxide precipitation or A-Level Chemistry 9476 aqueous ion questions, this is a particularly useful case because it connects oxidation state, coordinate bonding and Brønsted–Lowry acid behaviour.
I ask students to compare two ions: one Fe²⁺ and one Fe³⁺. They belong to the same element. They can both bind water, often represented with six water molecules around the central ion. Why should one behave as a stronger acid than the other? The answer is that the positive charge is not equally distributed over an equally sized centre. The more highly charged iron(III) centre draws more electron density from the oxygen atoms of coordinated water, changing how readily those water ligands lose H⁺.
This is deeper reasoning based on the H2 Chemistry foundations of aqueous acid–base behaviour, metal oxidation states, coordination and practical ion identification. The detailed numerical hydrolysis model is an educational extension, not a claim that every Fe aqua-ion pKa must be memorised for the 2027 9476 examination. Follow the specified syllabus, practical notes and question data for the exact cohort.
At eduKate, our small-group model of no more than three students aims to identify the first unstable idea, then check it on a changed example. Parents searching from Hougang should confirm actual teaching location, lesson timetable and availability rather than inferring a physical Hougang classroom from this article.
Hougang Chemistry Tutor Library · Official SEAB H2 Chemistry syllabus listing
Begin with the two iron oxidation states
Iron has atomic number 26. It commonly forms ions in oxidation states +2 and +3.
The conventional ion symbols are:
Iron(II): Fe²⁺.
Iron(III): Fe³⁺.
In many aqueous situations, water molecules coordinate directly to these metal centres.
The familiar simplified coordination formulae are:
[Fe(H₂O)₆]²⁺.
[Fe(H₂O)₆]³⁺.
Each contains a central iron ion surrounded by six water ligands in an approximately octahedral coordination environment.
The H₂O ligands are formally neutral, so the charge of each overall complex matches the oxidation state of the central iron in these expressions.
That difference in charge is the first important clue.
The two complexes do not behave identically merely because both include six water molecules.
Why coordinated water matters
Water contains polar O–H bonds and oxygen lone pairs.
A water molecule can donate a lone pair from oxygen to a metal centre, forming a coordinate interaction.
In the simplified aqueous iron complex, oxygen atoms of water molecules coordinate to the iron centre. The interaction affects electron density in the water ligands.
A strongly positively charged metal centre attracts electron density strongly.
That interaction can make an O–H bond in a coordinated water molecule more polarised and can increase the tendency for its hydrogen to be donated as a proton to another water molecule.
So the complex ion itself can behave as a Brønsted acid, even though it is not an ordinary molecule written as HCl or HNO₃.
The acidic proton comes from a water ligand coordinated to the metal, not directly from a bare iron ion containing a hydrogen atom.
This distinction helps explain why the solution pH can depend on the dissolved metal ion.
Write the first iron(III) hydrolysis equilibrium
For the hydrated iron(III) ion:
[Fe(H₂O)₆]³⁺ + H₂O ⇌ [Fe(H₂O)₅(OH)]²⁺ + H₃O⁺.
What changed?
One water molecule originally coordinated to iron has lost a proton, becoming a coordinated OH⁻ group.
A surrounding water molecule accepts that proton, becoming H₃O⁺.
The overall complex charge changes from +3 to +2 because one neutral H₂O ligand has become a negatively charged OH ligand.
But the iron atom has not changed its oxidation state.
The charge-balance check for the hydrolysed complex is:
Fe oxidation state + five neutral water ligands + one OH⁻ ligand = complex charge +2.
So:
x + 5(0) − 1 = +2.
Therefore:
x = +3.
Iron remains iron(III) throughout this acid–base process.
This is an important distinction between hydrolysis and reduction.
Write the analogous iron(II) hydrolysis equilibrium
A corresponding first hydrolysis of hydrated iron(II) can be written as:
[Fe(H₂O)₆]²⁺ + H₂O ⇌ [Fe(H₂O)₅(OH)]⁺ + H₃O⁺.
The process is structurally similar.
One coordinated water ligand donates H⁺. The resulting complex carries one less positive charge, because one water ligand has become a hydroxide ligand.
The central iron remains oxidation state +2:
x − 1 = +1, giving x = +2.
Both Fe²⁺ and Fe³⁺ aqua complexes can behave as acids through this mechanism.
The difference is the extent to which the equilibrium favours proton transfer.
Under comparable chemical conditions, iron(III) aqua complexes are considerably more acidic than their iron(II) counterparts.
A student should not write that Fe²⁺ can never donate a proton. The correct comparison is relative acid strength, not an absolute on/off rule.
Why Fe³⁺ polarises the O–H bond more strongly
Iron(III) carries charge +3, while iron(II) carries charge +2.
The higher oxidation state also corresponds to a more strongly contracting effective cation environment under the familiar ionic-size comparison.
Consequently Fe³⁺ has higher positive charge density and stronger polarising power toward the coordinated water molecules.
That draws electron density from oxygen more strongly, affecting the O–H bond electron distribution and making release of a proton more favourable.
The conceptual chain is:
Greater positive charge density → stronger polarisation of coordinating water → greater acidity of the coordinated water ligand → more favourable H₃O⁺ formation.
This is not a direct equation allowing you to calculate pH by simply dividing the ion charge by the ionic radius.
It explains the trend. Exact pH still depends on concentration, hydrolysis constants, other ligands, counterions and subsequent equilibria.
That is why a well-written H2 response gives a structural reason rather than an invented numerical pH.
What does “more acidic” mean quantitatively?
For an acid dissociation equilibrium, the acid dissociation constant Ka measures the equilibrium ratio of products and reactants under a specified convention.
A larger Ka generally indicates more extensive dissociation and stronger acid behaviour for a comparable pair under the relevant conditions.
A smaller pKa corresponds to a stronger acid, since:
pKa = −log₁₀ Ka.
Representative chemical references give the first hydrolysis of the iron(III) aqua ion a pKa of order 2–3, while the iron(II) hexaaqua ion has a much higher pKa, of order 9–10, depending on definition and solution environment.
These are broad reference ranges, not universal constants for all iron salts in every solvent.
Aqueous iron(III) can undergo multiple hydrolysis steps, ion pairing, polymerisation and precipitation. Its exact apparent acidity therefore depends on the defined chemical system.
The scientifically useful conclusion is that Fe³⁺-coordinated water is much more acidic than Fe²⁺-coordinated water.
Why Fe³⁺ solution can be acidic
When an iron(III) salt dissolves, hydrated Fe(III) species may form and undergo acid dissociation.
The hydrolysis equation generates hydronium:
[Fe(H₂O)₆]³⁺ + H₂O ⇌ [Fe(H₂O)₅(OH)]²⁺ + H₃O⁺.
A higher hydronium activity corresponds to a lower pH.
Thus a solution of an iron(III) salt, under compatible conditions, may be appreciably acidic even when the counterion is not itself a strong proton donor.
However, real Fe(III) salts can also be supplied in acidified solutions for stability. If the sample contains added HCl or another acid, its pH is influenced by that added acid as well as the iron hydrolysis equilibria.
Do not claim every measured low pH in an FeCl₃ bottle is proof of hydrolysis alone.
A controlled comparison would need to specify the iron concentration, added acid and other ions.
Why iron(II) is less acidic but not necessarily perfectly neutral
Hydrated Fe²⁺ is less strongly polarising than Fe³⁺.
Its water-ligand proton-transfer equilibrium is correspondingly less favourable.
Under suitable comparable aqueous conditions, an Fe²⁺ solution can be only mildly acidic relative to a Fe³⁺ solution.
But “less acidic than Fe³⁺” does not mean “always exactly neutral at pH 7”.
The actual pH can depend on dissolved oxygen, counterions, added acids, concentrations, temperature and oxidation during handling.
Iron(II) can be oxidised in air under suitable conditions, generating Fe(III) species that may then hydrolyse or precipitate.
So a real mixture left standing in air may change chemical composition over time.
An exam question that asks about intrinsic hydrated-ion acidity usually assumes the specified metal oxidation states remain controlled.
Keep acid–base hydrolysis separate from iron redox
There are two different kinds of transformations to distinguish.
Hydrolysis:
[Fe(H₂O)₆]³⁺ + H₂O ⇌ [Fe(H₂O)₅(OH)]²⁺ + H₃O⁺.
The Fe oxidation state is still +3.
Redox:
Fe³⁺ + e⁻ → Fe²⁺.
The Fe oxidation state changes from +3 to +2 because the metal centre gains an electron.
Both transformations can alter the species present in an aqueous sample, but they answer different questions.
One is proton transfer from coordinated water.
The other is electron transfer changing iron’s oxidation state.
A student who says the hydrolysed [Fe(H₂O)₅(OH)]²⁺ species contains Fe²⁺ has mistaken the overall complex charge for the metal oxidation state.
The correct charge assignment includes the negatively charged hydroxide ligand.
What happens when hydroxide is added to Fe²⁺?
Under the familiar school qualitative-analysis conditions, Fe²⁺ reacts with hydroxide ions to give a green precipitate represented as iron(II) hydroxide:
Fe²⁺ + 2OH⁻ → Fe(OH)₂(s).
The hydrated-ion model can represent this more explicitly as deprotonation of coordinated waters, forming an insoluble iron(II) hydroxide-containing phase.
The simple net ionic equation remains useful for standard exam stoichiometry.
The solid is commonly described as green when freshly prepared under suitable conditions.
The colour and its changes are evidence about the composition and oxidation state of the precipitating species.
But do not confuse precipitation with the initial first hydrolysis equilibrium. They are related stages of aqueous iron chemistry, with different amounts of proton removal and different resulting phases.
What happens when hydroxide is added to Fe³⁺?
A corresponding simplified Fe³⁺ precipitation equation is:
Fe³⁺ + 3OH⁻ → Fe(OH)₃(s).
The familiar iron(III) hydroxide precipitate is red-brown, often discussed as a hydrated iron(III) oxide or hydroxide-like solid rather than a perfectly pure molecular Fe(OH)₃ phase in every actual sample.
Three hydroxide equivalents are required per Fe³⁺ in the simple net ionic expression.
The iron(III) aqua complex can undergo multiple deprotonation steps before a neutral, sparingly soluble hydroxide-containing phase forms.
That is why the hydrated-ion acid concept connects naturally to qualitative analysis.
The iron centre still has oxidation state +3 throughout the direct hydroxide-precipitation stoichiometry.
No electron transfer is needed simply to form iron(III) hydroxide from Fe³⁺ and OH⁻.
Compare Fe²⁺ and Fe³⁺ side by side
| Property | Iron(II) system | Iron(III) system |
|---|---|---|
| Simple aqua complex | [Fe(H₂O)₆]²⁺ | [Fe(H₂O)₆]³⁺ |
| Metal oxidation state | +2 | +3 |
| Formal complex charge | +2 | +3 |
| Polarising power | Weaker in comparison | Stronger |
| Relative aqua-ion acidity | Weaker | Stronger |
| Typical first hydrolysis product | [Fe(H₂O)₅(OH)]⁺ | [Fe(H₂O)₅(OH)]²⁺ |
| Simple hydroxide precipitate | Fe(OH)₂, green | Fe(OH)₃, red-brown |
| NaOH equivalents in simple net precipitation | 2 per Fe²⁺ | 3 per Fe³⁺ |
The table summarises the standard models. Exact colours, species distributions and apparent acidities depend on solution conditions and handling.
Why green iron(II) hydroxide can turn brown in air
A freshly formed iron(II) hydroxide precipitate can be oxidised by atmospheric oxygen.
A simplified balanced overall equation is:
4Fe(OH)₂(s) + O₂(g) + 2H₂O(l) → 4Fe(OH)₃(s).
The iron centres change from oxidation state +2 to +3.
The resulting iron(III) hydroxide-like material has a red-brown appearance in the standard school description.
This does involve redox because electrons are transferred to oxygen.
By contrast, the original precipitation:
Fe²⁺ + 2OH⁻ → Fe(OH)₂(s)
does not change iron’s oxidation state.
The two processes can occur one after another in a practical sample. That is why an observation taken immediately and another taken after standing in air may be different.
A student who sees a brown precipitate after a delay should consider whether a green Fe(II) precipitate was oxidised, instead of assuming the sample contained only Fe³⁺ from the beginning.
Why the real solution colour can be more complicated
Aqueous Fe(II) is often described as pale green under ordinary conditions.
Aqueous Fe(III) may appear pale yellow, yellow or orange depending on the hydrated and hydrolysed species, their concentration and other ligands.
In fact, the isolated hexaaquairon(III) species itself has a different characteristic appearance from some of the more strongly coloured hydrolysed species.
The familiar yellow-orange colour of many Fe(III) solutions can therefore involve hydrolysis products and their electronic absorption, not simply one unchanging Fe³⁺ aqua formula.
This is why a colour label should not be treated as an exact molecular-composition measurement.
A practical qualitative-analysis scheme combines reagent response, precipitate behaviour and other evidence, rather than identifying a complete iron coordination environment from one shade.
For a more detailed treatment, see Chemistry LibreTexts: aqueous iron complexes and hydroxide reactions.
A hypothetical Ka example comparing two metal aqua acids
To practise the mathematics, imagine two hypothetical monoprotonic aqua complexes called A³⁺ and B²⁺. These are not actual measured iron equilibrium constants.
Suppose the initial concentration of each complex is:
C = 0.0100 mol dm⁻³.
Give their first hydrolysis constants as:
Ka(A³⁺) = 1.0 × 10⁻³.
Ka(B²⁺) = 1.0 × 10⁻⁹.
These values are deliberately chosen so A³⁺ is much more acidic than B²⁺ in this simplified model.
Assume neither undergoes additional hydrolysis or precipitation, and ignore other acid–base species.
The problem asks how their equilibrium hydronium concentrations compare.
Solve the stronger aqua acid without an invalid approximation
For A³⁺:
A³⁺ + H₂O ⇌ AOH²⁺ + H⁺.
Let the dissociated amount be x.
Then:
Ka = x²/(0.0100 − x).
So:
1.0 × 10⁻³ = x²/(0.0100 − x).
Rearrange:
x² + 0.00100x − 0.0000100 = 0.
Use the positive quadratic root:
x = [−0.00100 + √((0.00100)² + 4(0.0000100))]/2.
The result is approximately:
x = 0.00270 mol dm⁻³.
Therefore:
[H⁺] ≈ 0.00270 M.
pH ≈ −log₁₀(0.00270) ≈ 2.57.
The dissociation is about 27% of the initial concentration, so the small-x approximation would have been poor. Solving the quadratic is appropriate for these hypothetical values.
The result illustrates how a relatively strong aqua acid can generate appreciable hydronium.
Solve the weaker aqua acid
For B²⁺, use the much smaller hypothetical Ka:
Ka = 1.0 × 10⁻⁹.
At the same C = 0.0100 M, the small-dissociation approximation is appropriate:
[H⁺] ≈ √(KaC).
So:
[H⁺] ≈ √[(1.0 × 10⁻⁹)(0.0100)].
[H⁺] ≈ √(1.0 × 10⁻¹¹) ≈ 3.16 × 10⁻⁶ M.
Then:
pH ≈ −log₁₀(3.16 × 10⁻⁶) ≈ 5.50.
Under the chosen hypothetical model, the weaker aqua ion solution has much higher pH.
But these pH values do not claim that 0.0100 M aqueous Fe(III) and Fe(II) necessarily have exactly pH 2.57 and 5.50.
Real iron hydrolysis involves additional species, counterion effects, activity corrections and possible precipitation or oxidation. The numerical example teaches the method, while the actual iron chemistry provides the qualitative explanation.
Why exact Fe(III) solution pH is not one fixed number
Iron(III) hydrolysis is not restricted to one simple dissociation step at every pH.
After the first hydrolysis, additional water ligands may lose protons. Hydrolysed iron species may associate into larger structures, and insoluble iron(III) hydroxide or oxide-related phases may form under suitable conditions.
The distribution depends on total iron concentration, pH, counterions, temperature and solution history.
Therefore a pH calculation treating Fe³⁺ as a single ideal weak monoprotic acid may be a useful teaching approximation but not a complete chemical speciation model.
The student should recognise when the problem supplies an explicit one-step model and when more information would be needed for a real sample.
The key qualitative truth—that higher-charged hydrated Fe³⁺ is more acidic than Fe²⁺ under suitable comparison conditions—does not require pretending one simple equation describes every possible iron solution.
How added chloride or other ligands can complicate the comparison
Real FeCl₂ and FeCl₃ solutions can contain chloride-related species.
Chloride ions may interact with iron centres through coordination equilibria depending on concentration and the chemical environment.
A solution prepared using an added strong acid, such as HCl, also contains hydronium from that acid.
Consequently comparing two commercial iron salts by measured pH alone is not always a controlled comparison of the intrinsic aqua-ion acid strengths.
If the scientific question is about why Fe³⁺-coordinated water is more acidic, we should conceptually compare the hydrated ions at otherwise comparable conditions.
If the question is about two actual prepared solutions, we must consider all added reagents and possible species.
This is the general laboratory principle of controlling other variables.
Can an ion be an acid and an oxidising agent at once?
Yes. An iron(III)-containing solution can participate in different reactions depending on the partner.
In the hydrolysis reaction:
[Fe(H₂O)₆]³⁺ + H₂O ⇌ [Fe(H₂O)₅(OH)]²⁺ + H₃O⁺,
the hydrated complex acts as an acid, donating a proton from a coordinated water ligand.
In the redox reaction:
Fe³⁺ + e⁻ → Fe²⁺,
Fe³⁺ acts as an oxidising agent, accepting an electron.
Those are different chemical roles.
The word acidic is not synonymous with oxidising, and having a positive ion charge does not automatically specify one reaction direction with every possible partner.
This is especially useful in integrated H2 questions where solution pH and redox potential are both discussed.
A good student labels the actual particle and transferred species: proton in acid–base chemistry, electron in redox chemistry.
A worked mole calculation with iron hydroxides
Suppose a purely theoretical qualitative-analysis question contains 0.00500 mol Fe²⁺, and sufficient OH⁻ is present to precipitate all iron as Fe(OH)₂ under the ideal net equation:
Fe²⁺ + 2OH⁻ → Fe(OH)₂(s).
The stoichiometric amount of OH⁻ required is:
n(OH⁻) = 2 × 0.00500 = 0.0100 mol.
Now consider an equally sized separate sample containing 0.00500 mol Fe³⁺, precipitated under:
Fe³⁺ + 3OH⁻ → Fe(OH)₃(s).
The stoichiometric amount of OH⁻ is:
n(OH⁻) = 3 × 0.00500 = 0.0150 mol.
The difference comes from the charges and hydroxide coefficients in the balanced net precipitation reactions.
It does not directly measure the first hydrolysis pKa of the aqua ions. Stoichiometric complete precipitation and the equilibrium acidity of initial dissolved complexes are different quantitative questions.
This distinction prevents students from using a 3:2 hydroxide mole ratio to infer that Fe³⁺ is exactly 1.5 times as strong an acid.
The common mistakes worth diagnosing
| Student statement | First wrong assumption | Repair question |
|---|---|---|
| Fe³⁺ is more acidic because it already contains more H atoms | Ignored coordinated water and proton origin | Which ligand loses H⁺? |
| Fe²⁺ aqua ions cannot be acidic | Confused weaker acidity with no acidity | Can a coordinated Fe²⁺ water ligand also deprotonate? |
| Hydrolysis of [Fe(H₂O)₆]³⁺ forms Fe²⁺ | Confused overall complex charge with metal oxidation state | What charge does the OH ligand contribute? |
| Adding OH⁻ is the same as adding an electron | Confused proton transfer and redox | Does the iron oxidation number change? |
| Every Fe³⁺ solution has one exact pH | Ignored concentration and complex speciation | What hydrolysis constants and other species are present? |
| The green Fe(OH)₂ precipitate can never change colour | Ignored oxidation by atmospheric oxygen | What happens to Fe(II) in air? |
| Brown iron(III) precipitate proves the solution was originally Fe³⁺ only | Ignored possible oxidation after sampling | Was a green intermediate observed earlier? |
| FeCl₃ acidity comes solely from chloride | Ignored hydrated cation acidity and other reagents | Why does NaCl with Cl⁻ behave differently? |
Seven independent practice questions
- Write the simplified hexaaqua formula for Fe²⁺ and Fe³⁺ in water.
- State why Fe³⁺ has greater polarising power toward coordinated water than Fe²⁺.
- Write the first hydrolysis equilibrium of [Fe(H₂O)₆]³⁺, including the hydronium product.
- Calculate the oxidation state of iron in [Fe(H₂O)₅(OH)]²⁺.
- Write the familiar Fe²⁺ and Fe³⁺ hydroxide precipitation equations.
- Explain why freshly green Fe(OH)₂ may turn red-brown on standing in air.
- Does the difference in complete-precipitation OH⁻ stoichiometry of Fe²⁺ and Fe³⁺ directly equal the ratio of their first hydrolysis constants?
Answers
1. [Fe(H₂O)₆]²⁺ and [Fe(H₂O)₆]³⁺. 2. Fe³⁺ has higher positive charge density and more strongly polarises the coordinated O–H bonds. 3. [Fe(H₂O)₆]³⁺ + H₂O ⇌ [Fe(H₂O)₅(OH)]²⁺ + H₃O⁺. 4. +3. 5. Fe²⁺ + 2OH⁻ → Fe(OH)₂(s); Fe³⁺ + 3OH⁻ → Fe(OH)₃(s). 6. Fe(II) can be oxidised by atmospheric O₂ to Fe(III), producing red-brown hydroxide-like material. 7. No. Precipitation stoichiometry and acid-dissociation equilibrium constants are different quantities.
If a student gets Question 4 wrong, the first repair is understanding ligand charge and oxidation state, not revisiting every transition-metal colour observation.
A seven-session JC1–JC2 learning route
- Session 1: compare Fe²⁺ and Fe³⁺ oxidation states and electron configurations.
- Session 2: draw the two hexaaqua complexes and identify the donor oxygen atoms.
- Session 3: explain how metal charge density affects coordinated water’s O–H bonds.
- Session 4: balance the first hydrolysis equilibrium for both ions.
- Session 5: solve one explicitly hypothetical Ka/pH comparison and check whether an approximation is justified.
- Session 6: connect the Fe²⁺ and Fe³⁺ hydroxide precipitation equations to familiar qualitative observations.
- Session 7: distinguish hydrolysis, precipitation and oxidation in one unfamiliar multi-stage aqueous iron question.
The final session is the most useful. It tests whether the learner can decide which chemical process a new observation represents.
What parents looking for H2 Chemistry tuition near Hougang can ask
Bring a marked transition-metal or aqueous ion question. Ask whether the student’s earliest error is confusing metal ion charge, coordinated water, hydrolysis, metal oxidation state or a later precipitation reaction.
An effective small-group or private session should finish with a changed metal aqua ion and an explanation of its expected acidity, without copying the iron example word for word.
At eduKate, the maximum three-student tutorial model is intended to preserve individual feedback. Confirm the exact lesson venue, current H2 syllabus experience, timetable and travel requirements before enrolling; a Hougang-focused Chemistry article does not prove that a physical class operates inside Hougang.
Why the concept matters beyond A Levels
Hydrated iron chemistry is important in environmental water systems, corrosion, analytical chemistry and materials processing.
The speciation of a metal depends on its oxidation state, ligands, pH and surrounding environment. These properties determine whether a metal remains dissolved, releases protons, forms a solid phase or undergoes electron transfer.
The lasting scientific habit is to distinguish proton transfer, electron transfer and precipitation before attempting to explain an observed colour or a measured pH.
That skill supports advanced Chemistry and environmental or chemical engineering, regardless of future university choices.
Frequently asked questions
Why is aqueous Fe³⁺ more acidic than Fe²⁺?
The higher charge density of Fe³⁺ more strongly polarises the O–H bonds of coordinated water, making proton transfer to the solvent more favourable.
Do iron ions themselves contain acidic hydrogen?
The protons originate from water ligands coordinated to the metal centre, not from a bare Fe ion.
What is the first hydrolysis equation for Fe³⁺?
[Fe(H₂O)₆]³⁺ + H₂O ⇌ [Fe(H₂O)₅(OH)]²⁺ + H₃O⁺.
Does the iron oxidation state change during hydrolysis?
No. A coordinated water becomes hydroxide, changing overall complex charge while the central metal remains Fe(III) or Fe(II).
What is the usual Fe²⁺ hydroxide precipitate colour?
Fresh Fe(OH)₂ is conventionally described as green under appropriate conditions.
What is the usual Fe³⁺ hydroxide precipitate colour?
Red-brown.
Why can a green Fe²⁺ precipitate become brown?
Atmospheric oxygen can oxidise Fe(II) to Fe(III), giving red-brown hydroxide-like products.
Is detailed Fe aqua-ion pH calculation compulsory for 9476?
The detailed Ka model here is enrichment to connect prescribed acid–base and coordination principles. Use the official syllabus and values supplied in the examination question to determine required depth.
Continue learning and references
Why AlCl₃ is acidic but NaCl is nearly neutral · Why Fe³⁺ oxidises iodide to iodine · H2 Paper 4 qualitative analysis of metal ions · Why green iron(II) hydroxide turns brown in air · Chemistry LibreTexts: iron aqua ions and hydroxide reactions · Chemistry LibreTexts: acidity of hexaaqua ions · Hougang Chemistry Tutor Library · eduKate Singapore JC Chemistry enquiries.
Scientific-model note: The numerical Ka example uses invented constants for illustrative metal aqua ions, not measured Fe²⁺ or Fe³⁺ pH values. Real iron solutions can involve multiple hydrolysed species, oxidation, ion pairing and precipitation. Practical work requires approved supervised laboratory conditions.