Chemistry 11–12 · Year 12

Measuring the equilibrium constant of the iron(III) thiocyanate reaction by colourimetry

Module 5: Equilibrium and Acid Reactions

Practical, model not builtMedium risk

School laboratory, not for home

In a school laboratory, with a teacher supervising, under the school's risk assessment. Not for home.

This site has no interactive model of its own. Where a step or a material names a Concept Studio model, simulation or tool, it has not been built; an external simulation a step names (for example PhET) is not part of this site.

The idea

The concentration of the coloured complex at equilibrium, read from a calibration curve, fixes the equilibrium constant for a set of mixtures whatever their starting ratio.

Safety card

Medium riskLearners carry it out, with a teacher supervising

Setting: In a school laboratory, with a teacher supervising, under the school's risk assessment. Not for home.

Hazards

0.1 mol/L nitric acid irritant
iron(III) nitrate irritant
thiocyanate harmful; keep away from strong acid heating

Controls

pipette fillers
eye protection
no heating of thiocyanate mixtures
wastes to a labelled container

Note

NSW Department of Education Chemical Safety in Schools (CSIS) package, 2021 Technical Update: Section 1.7 (risk assessment) and Volume 2 Appendix D (generic assessment advice and DoE chemical categories); record a RiskAssess (riskassess.com.au) risk assessment before the lesson and check the school's hazardous chemical register (CSIS Section 1.9) for local restrictions.

What you need

0.200 mol/L iron(III) nitrate in 0.1 mol/L nitric acid, 50 mL (for calibration, where a large excess of Fe3+ converts all SCN- to complex)
0.00200 mol/L iron(III) nitrate in 0.1 mol/L nitric acid, 50 mL, and 0.00200 mol/L potassium thiocyanate, 50 mL
0.1 mol/L nitric acid as diluent, 100 mL
Colourimeter with a blue-green filter (about 470 nm) or a smartphone colourimetry app with a fixed light box, cuvettes
10.00 mL and 5.00 mL bulb pipettes; two 50 mL burettes, one for the 0.00200 mol/L thiocyanate (the 1.00 to 5.00 mL portions) and one for the acid that makes each tube up to 10.0 mL; 25 mL volumetric flasks, five; test tubes

How to do it

Calibration: make five standards by adding 1.00, 2.00, 3.00, 4.00 and 5.00 mL of 0.00200 mol/L thiocyanate to 10.0 mL of 0.200 mol/L iron(III) nitrate and diluting each to 25.0 mL with the acid; in these the complex concentration equals the thiocyanate added (excess iron drives the reaction to completion).
Measure the absorbance of each standard at 470 nm against an acid blank and plot absorbance against complex concentration; fit the straight line through the origin.
Equilibrium mixtures: in five tubes mix 5.00 mL of 0.00200 mol/L iron(III) nitrate (bulb pipette) with 1.00, 2.00, 3.00, 4.00 and 5.00 mL of 0.00200 mol/L thiocyanate, diluting each to 10.0 mL with acid from the burette; measure each absorbance.
Read the equilibrium complex concentration from the calibration line; subtract it from the initial concentrations of each ion to get their equilibrium values; calculate K for each mixture and report the mean and spread.
Enter the same volumes in the simulation and compare the predicted absorbances and K with the measured set.

What you should see

The calibration line is straight (absorbance rises in proportion to complex concentration across the standards, 8.0e-5 to 4.0e-4 mol/L). The calibration assumes the 0.080 mol/L excess of iron(III) converts nearly all the thiocyanate to complex: for K = 100 L/mol that is 89 percent and for K = 200 L/mol 94 percent, a bias the learner should state. For the mixture of 5.00 mL of each 0.00200 mol/L solution (each ion 0.00100 mol/L at the start), K = 100 L/mol predicts 0.0839 mmol/L of complex (8.4 percent conversion) and K = 200 L/mol 0.1459 mmol/L. The test of the method is that the five mixtures give the same K within their measurement spread; published values depend on ionic strength and temperature, so the class compares its mixtures with each other rather than with one textbook figure.

What changes

What you change: starting ratio of iron(III) to thiocyanate
What you measure: equilibrium concentration of [FeSCN]2+ (from absorbance) and the calculated K
What you keep the same: temperature
acid concentration (ionic strength)
wavelength and cuvette
total volume

Common misconceptions

Each of these ideas is wrong, and the activity is a chance to test it.

The equilibrium constant changes when the starting amounts change (K stays the same; the equilibrium concentrations change).
Absorbance measures the amount of iron present (it measures the coloured complex only; free Fe3+ is nearly colourless at this wavelength).
A bigger K means a faster reaction (K describes the position, not the rate).

Curriculum references

The NSW syllabus outcomes and Australian Curriculum v9 codes this activity supports. They are references, not a verified or complete curriculum alignment.

Chemistry Stage 6 Syllabus (2017), NESA; the current syllabus, taught in 2026 (codes read from the syllabus document)CH12-12 CH11/12-4 CH11/12-5 CH11/12-6
Chemistry 11-12 Syllabus (2025), NESA; implemented from 2028, not yet taughtCH-12-01 CH-12WS-04 CH-12WS-05 CH-12WS-06
Australian Curriculum v9No Australian Curriculum v9 code is listed.

Sources

The pages the author read to write this activity.

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