Colorimetry quantitative analysis determining formula of transition metal complex ion colorimeter measurements calibration graph GCE AS A2 IB inorganic chemistry revision notes

Periodic Table - Transition Metal Chemistry - Doc Brown's Chemistry Revising Advanced Level Inorganic Chemistry Periodic Table Revision Notes

Appendix 9 Colorimetry – quantitative analysis and determining the formula of a complex ion

Using a colorimeter to measure the concentration of transition metals in solution and a method of determining the formula of a complex ion.

GCSE/IGCSE Periodic Table Revision Notes * GCSE/IGCSE Transition Metals Revision Notes

Doc Brown's Chemistry Advanced Level Pre-University Chemistry Revision Study Notes for UK IB KS5 A/AS GCE advanced level inorganic chemistry students US K12 grade 11 grade 12 inorganic chemistry - 3d block transition metal chemistry Sc Ti V Cr Mn Fe Co Ni Cu Zn

INORGANIC Part 10 3d block TRANSITION METALS sub–index: 10.1–10.2 Introduction 3d–block Transition Metals * 10.3 Scandium * 10.4 Titanium * 10.5 Vanadium * 10.6 Chromium * 10.7 Manganese * 10.8 Iron * 10.9 Cobalt * 10.10 Nickel * 10.11 Copper * 10.12 Zinc * 10.13 Other Transition Metals e.g. Ag and Pt * Appendix 1. Hydrated salts, acidity of hexa–aqua ions * Appendix 2. Complexes & ligands * Appendix 3. Complexes and isomerism * Appendix 4. Electron configuration & colour theory * Appendix 5. Redox equations, feasibility, E^ø * Appendix 6. Catalysis * Appendix 7. Redox equations * Appendix 8. Stability Constants and entropy changes * Appendix 9. Colorimetric analysis and complex ion formula * Appendix 10 3d block – extended data * Appendix 11 Some 3d–block compounds, complexes, oxidation states & electrode potentials * Appendix 12 Hydroxide complex precipitate 'pictures', formulae and equations

Advanced Level Inorganic Chemistry Periodic Table Index: Part 1 Periodic Table history * Part 2 Electron configurations, spectroscopy, hydrogen spectrum, ionisation energies * Part 3 Period 1 survey H to He * Part 4 Period 2 survey Li to Ne * Part 5 Period 3 survey Na to Ar * Part 6 Period 4 survey K to Kr and important trends down a group * Part 7 s–block Groups 1/2 Alkali Metals/Alkaline Earth Metals * Part 8 p–block Groups 3/13 to 0/18 * Part 9 Group 7/17 The Halogens * Part 10 3d block elements & Transition Metal Series * Part 11 Group & Series data & periodicity plots * All 11 Parts have their own sub–indexes near the top of the pages

Appendix 9. Colorimetry – quantitative analysis and determining the formula of a complex ion

(a) What is colorimetry? what is a colorimeter? how does a colorimeter work?

Colorimetry is method of determining the concentration of a substance by measuring the relative absorption of light (usually visible) with respect to a known concentration of the substance.
- For theory see Electron configuration & colour theory
- and Colour and uv-visible absorption spectroscopy
The instrument by which these measurements are made is called a colorimeter (illustrated below).
Light from a suitable source is passed through a light filter to select the most appropriate wavelength of light, some of which is then absorbed by the solution held in a special glass cuvet (a sort of 'test tube').
The amount of light absorbed is called, and measured as, the absorbance which is a function of the coloured solute concentration.
Most expensive instruments use a double beam system of two cuvets, one is a 'blank' of water and one the actual coloured solution under test, two photocells and sophisticated optics of lenses and mirrors which need not concern as at all.
Cheaper colorimeters (i.e. in school and illustrated above) allow you to put in a cuvet of 'colourless' water, zero the instrument i.e. set it to read zero absorbance, replace with a cuvet of the coloured solution and simply read of the 'absorbance'. The 'zeroing' is necessary because even the apparently 'colourless blank' of glass cuvet and water can absorbed a tiny amount of light. This procedure eliminates this error.
The filter is chosen to select the band of wavelengths which are most strongly absorbed by the coloured solution e.g. this is illustrated on the diagram above, and in the table below, by using a yellow filter to use in measuring the concentration of a blue coloured solution like copper(II) sulfate or its ammine/amine complex.

The wavelength (nm) of the observed transmitted colour of the solution	*The observed transmitted colour of the solution ( as in the diagram above)**	The complementary colour of the solution i.e. the colour of the filter
400–435	violet	yellowish–green
435–480 *	blue *	yellow *
480–490	greenish–blue	orange
490–500	bluish–green	red
500–560	green	purple
560–580	yellowish–green	violet
580–595	yellow	blue
595–610	orange	greenish–blue
610–750	red	bluish–green

Although the table illustrates the 'complementary' colour relationship between the solution and the filter, in practice it is better to try several filters on a typical concentration of the solution under test to see which filter gives the highest absorption value i.e. gives you maximum sensitivity and hence maximum accuracy in your measurements.

How do we use colorimetry to measure the concentration of a transition metal ion?
- If an aqueous transition metal ion is intensely coloured, its concentration can be measured directly e.g. manganese concentration can be measured if it is oxidised to the deep purple manganate(VII) ion, MnO₄^–.
- However, many ions are not as intense as the MnO₄^– ion, BUT if a suitable ligand or complexing agent is added, then a more intensely coloured complex may be formed, from which accurate measurements of concentration can be made e.g.
  - The blue hexaaquacopper(II) ion forms a deeper violet–blue ammine complex with ammonia.
  - Yellow–brown iron(III) ions form a deep blood–red complex with the thiocyanate ion (SCN^–) by mixing it with ammonium/potassium cyanate.
- Once the method of producing a more intense colour is established, you then need to derive a calibration curve.
  - This is done by measuring the absorbance of solutions of known concentrations of the coloured complex and plotting the calibration curve/graph (see right of colorimeter diagram).
  - The known concentration range should include any likely absorbance measured from the solutions under test i.e. those whose concentration is being determined.
  - Generally at low concentrations the calibration curve is linear i.e. it obeys Beer's Law (Beer–Lambert Law). Without going into the mathematics of Beer's Law and absorption, it basically states that a solution's absorbance is directly proportional to the concentration of the coloured solute.
    - For various reasons the calibration may curve upwards (positive deviation from Beer's Law) or curve over (negative deviation from Beer's Law). However, a linear or otherwise calibration curve still shows increasing absorption with increasing concentration and curved calibration graphs are acceptable, if not advisable, if your methodology is accurate.
How can we use colorimetry to deduce the formula of a complex?
- The method depends on measuring the absorbencies of solutions containing different ratios of transition metal ion to complexing agent.
- Just for the sake of argument, imagine that one mole of transition metal (M³⁺) ion reacts with two moles of a monodentate ligand (X^–). The reaction equation for the ligand displacement reaction to form the complex would be:
  - (i) [M(H₂O)₆]³⁺_(aq) + 2X^–_(aq) [M(H₂O)₄X₂]⁺_(aq) + 2H₂O_(l)
- There are two basic approaches as to how you vary the transition metal ion – ligand ratio and the results illustrated in the diagram below.
- Method (1) The mole ratio method keeping Mⁿ⁺ constant and gradually increasing the number of moles of ligand X from zero to a large molar excess.
  - From 0.0 to 2.0 moles of X^– added per 1.0 moles of M³⁺ there is a steady rise in absorbance as more and more of the complex is formed. From over 2.0 of X^– per mol M³⁺ there is a more gradual rise in absorbance.
    - The point of intersection of the two linear portions of the graph gives you the X^–/M³⁺ ratio in the complex i.e. 2.0
    - Note that ligand exchange reactions are equilibrium reactions so you don't go to the maximum absorbance at 2.0 but the change in the 'rate of change' of absorbance does give the ratio.
    - If you do this for the reaction between iron(III) ions and thiocyanate ions which gives a blood–red complex, you obtain the change in graph gradient at 1.0 mol of CNS^– to 1.0 mol of Fe³⁺ for the reaction ...
      - (ii) [Fe(H₂O)₆]³⁺_(aq) + CNS^–_(aq) [Fe(H₂O)₅CNS]²⁺_(aq) + H₂O_(l)
    - For dilute solutions of copper(II) ions and ammonia the graph gradient change occurs at 4.0 moles of NH₃ per mol of Cu²⁺ for the reaction ...
      - (iii) [Cu(H₂O)₆]³⁺_(aq) + 4NH_3(aq) [Cu(H₂O)₂(NH₃)₄]²⁺_(aq) + 4H₂O_(l)
- Method (2) The continuous variations method in which you start with zero moles of Mⁿ⁺ and an excess of the ligand X. In each successive mixture you then increase the amount of Mⁿ⁺ and decrease the amount of X^– and keep the total moles of Mⁿ⁺ and X^– constant.
  - For the sake of argument, if you assume both stock solutions are the same molarity, then the ratio of the volumes automatically gives you the X/M ratio in the mixture.
  - For the fictitious M³⁺ and F^– complex, the peak will occur between the 6X^– : 4M³⁺ mixture (ratio 1.5) and the 7X^– : 3M³⁺ mixture (ratio 2.33). Theoretically the peak should occur at a ratio of 2.0 for X^–/M³⁺ i.e. theoretically, in terms of a total volume of 10 units, it means a 6.66X^– : 3.33M³⁺ mixture by volume.
  - For the iron(III)–thiocyanate complex in reaction (ii) the peak will occur at 5CNS^–: 5Fe³⁺
    - i.e. a CNS^–/Fe³⁺ ratio of 1.0 for the complex [Fe(H₂O)₅CNS]²⁺
  - For the copper(II)–ammonia complex in reaction (iii) the peak will occur at 8NH₃ : 2Cu²⁺
    - i.e. an NH₃/Cu²⁺ ratio of 4.0 for the complex [Cu(NH₃)₄]²⁺ or [Cu(H₂O)₂(NH₃)₄]²⁺.
- In both cases, you work out the mole ratios present in the mixtures from the known concentrations and volumes of the stock solutions of Mⁿ⁺ and X^– used in each experiment.
OTHER APPLICATIONS of COLORIMETRY
- The manganate(VII) ion, MnO₄^–, e.g. in potassium manganate(VII) solution, is a brilliant purple colour and its concentration in very dilute solution can be measured by using colorimetry i.e. by comparing the absorbance of the solution versus a calibration graph of known concentrations of the manganate(VII) ion.
  - This can be applied to kinetic experiments e.g. measuring the rate of reaction of acidified potassium manganate(VII) when it oxidises ethanedioic acid/ethanedioate ion. The reaction is autocatalysed by the manganese(II) ions formed.
  - Therefore, theoretically, via the intensity of the manganate(VII) ion colour, you can investigate the effects of changing the concentration of potassium manganate(VII), ethanedioic acid, dilute sulfuric acid and manganese(II) sulfate.
  - A simplistic view would be to write an overall rate expression such as ...
  - rate = k[MnO₄^–]^a[(COOH)₂]^b[H⁺]^c[Mn²⁺]^d, where a, b, c and d represent the orders of reaction.
  - However:
    - (i) good results are not easy to obtain, specifically due to the effect of autocatalysis
    - (ii) it is unlikely the rate expression is as simple as you will encounter in your pre–university course!
    - See advanced level kinetic notes case study 2c.2 on this reaction
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Scandium * Titanium * Vanadium * Chromium * Manganese Iron * Cobalt * Nickel * Copper * Zinc * Silver & Platinum

Introduction 3d–block Transition Metals * Appendix 1. Hydrated salts, acidity of hexa–aqua ions * Appendix 2. Complexes & ligands * Appendix 3. Complexes and isomerism * Appendix 4. Electron configuration & colour theory * Appendix 5. Redox equations, feasibility, E^ø * Appendix 6. Catalysis * Appendix 7. Redox equations * Appendix 8. Stability Constants and entropy changes * Appendix 9. Colorimetric analysis and complex ion formula * Appendix 10 3d block – extended data * Appendix 11 Some 3d–block compounds, complexes, oxidation states & electrode potentials * Appendix 12 Hydroxide complex precipitate 'pictures', formulae and equations

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