In this experiment, the wavelengths of light absorbed by metal complexes are examined, and the relationships among the color of the complex, ligands, coordination structure, and electronic state are discussed.
In university inorganic chemistry and instrumental analysis experiments, ultraviolet-visible absorption spectra of copper, nickel, cobalt, iron, and other complexes may be measured to study the relationship between d-d transitions, ligand-field splitting, and color.
In a report on the absorption spectrum of a complex, it is not sufficient simply to write that “a peak appeared” or “the color was blue.”
It is important to explain around which wavelength absorption occurred, what electronic transition caused that absorption, and how the absorbed light corresponds to the observed color.
The type of ligand, oxidation state of the metal ion, coordination structure, concentration, contamination of the cell, and baseline correction can also be used as points for discussion.
This article clearly explains how to interpret absorption spectra of complexes, the relationship between d-d transitions and color, how to discuss absorption peaks, sources of error, points for improvement, and discussion examples that can be used in reports.
Note:
This article is a reference intended to assist with discussions of results obtained in chemistry experiments at universities and similar institutions.
For the actual measurement conditions, measurement wavelengths, solution preparation, handling of cells, operation of instruments, and safety precautions, always follow the instructions in your university’s laboratory manual and those given by your instructor or TA.
- What Is the Absorption Spectrum of a Complex?
- Results to Examine in an Absorption Spectrum
- Reference Experimental Values for Absorption Spectra of Complexes and Examples of Analysis of d-d Transitions and Color
- Reference Experimental Conditions
- Relationship Between Absorption Spectra and the Color of Complexes
- Examples of Absorption Maxima and Colors of Representative Complexes
- Calculation of Absorbance Using the Beer-Lambert Law
- Example Measurements of the Relationship Between Absorbance and Concentration
- Example Calculation of Molar Absorptivity
- Changes in Absorption Maxima Caused by Ligands
- Ligand-Field Strength and Absorption Wavelength
- Example of Determining Transition Energy From Absorption Wavelength
- Example of Changes in the Absorption Spectrum Caused by Complex Formation
- Examples When Absorbance Is Not Proportional to Concentration
- Example Organization of Complex Color and Absorption Spectra
- Example of Variation in Measured Values
- Example of How to Write the Results
- Points for Connecting the Results to the Discussion
- Example Discussion
- Summary
- What Is a d-d Transition?
- Relationship Between Ligand-Field Splitting and Absorption Wavelength
- Relationship Between the Color of a Complex and the Light It Absorbs
- Discussion of the Maximum Absorption Wavelength
- What Can Be Learned From the Magnitude of Absorbance?
- The Beer-Lambert Law and Complex Concentration
- Discussion When a Peak Is Broad
- Discussion When Multiple Peaks Are Present
- Difference From Charge-Transfer Transitions
- Spectral Changes When the Ligand Is Changed
- Spectral Changes When the Oxidation State Is Changed
- Relationship Between Coordination Structure and Absorption Spectrum
- Example Discussion of the Absorption Spectrum of a Copper Complex
- Example Discussion of the Absorption Spectrum of a Nickel Complex
- Example Discussion of the Absorption Spectrum of an Iron Complex
- Example Discussion of the Absorption Spectrum of a Cobalt Complex
- Errors When the Concentration Is Too High
- Errors When the Concentration Is Too Low
- Importance of Blank Correction
- Errors Caused by Cell Contamination or Bubbles
- Effect of the Solvent
- When the Complex Changes During Measurement
- Discussion When the Results Differ From Literature Values
- When the Result Can Be Considered Good
- Example Discussion When the Measurement Did Not Go Well
- How to Write Points for Improvement
- Difference Between a Superficial Discussion and a Good Discussion
- Examples of Expressions That Can Be Used in Reports
- Points to Check When Discussing the Absorption Spectrum of a Complex
- Summary
What Is the Absorption Spectrum of a Complex?
The absorption spectrum of a complex is a graph showing how much light of each wavelength is absorbed by a complex solution or complex sample.
Wavelength is commonly plotted on the horizontal axis and absorbance on the vertical axis, and when absorption becomes large at a particular wavelength, it is observed as a peak.
In transition-metal complexes, absorption may appear in the ultraviolet-visible region because of electronic transitions related to the d orbitals of the metal ion or charge-transfer transitions between the metal and ligands.
Complexes that absorb visible light show color corresponding to the light that is not absorbed or to its complementary color.
In a report, the positions of absorption peaks, magnitude of absorbance, peak shape, and sample color are compared and discussed.
Results to Examine in an Absorption Spectrum
In the absorption spectrum of a complex, check the wavelength of each peak, absorbance, peak width, presence or absence of shoulder peaks, and condition of the baseline.
It is also important to consider the relationship between the color of the measured complex and the wavelengths being absorbed.
Main Items to Include in the Results
- Type of complex measured
- Solvent
- Sample concentration
- Cell length
- Measurement wavelength range
- Maximum absorption wavelength
- Absorbance at the maximum absorption wavelength
- Number of peaks
- Peak width and shape
- Color of the sample solution
- Whether a blank measurement was performed
- Comparison with literature values or standard samples
Example of How to Write the Results:
When the absorption spectrum of the complex solution was measured, an absorption peak was observed in the visible region.
The maximum absorption wavelength was around ○○ nm, indicating that light at this wavelength was absorbed relatively strongly.
Because the sample solution appeared blue, the complementary color of the absorbed light is considered to have appeared as the observed color.
Reference Experimental Values for Absorption Spectra of Complexes and Examples of Analysis of d-d Transitions and Color
Here, reference experimental values are organized for measuring absorption spectra of metal complexes and discussing absorption maxima, d-d transitions, complex color, and ligand-field strength.
Absorbance, wavelength, molar absorptivity, complementary colors, types of ligands, concentration dependence, and measurement errors are summarized in a form that can be readily used in reports.
In transition-metal complexes, the d orbitals of the metal ion are split by the influence of ligands.
The phenomenon in which light is absorbed and an electron moves from a lower-energy d orbital to a higher-energy d orbital is called a d-d transition.
Because the complementary color of the light absorbed by a complex becomes the observed color, the absorption spectrum and color of a complex are closely related.
Reference Experimental Conditions
| Item | Details |
|---|---|
| Samples measured | Aqueous complex solutions of Cu(II), Ni(II), Co(II), Fe(III), Cr(III), and others |
| Measurement method | Ultraviolet-visible spectrophotometry |
| Measurement wavelength range | 350–800 nm |
| Cell length | 1.00 cm |
| Evaluation items | Absorption maxima, absorbance, molar absorptivity, d-d transitions, complementary colors, ligand-field strength |
| Main sources of error | Concentration preparation, blank correction, cell contamination, incomplete complex formation, pH changes, precipitation |
Relationship Between Absorption Spectra and the Color of Complexes
When a complex absorbs light of a particular wavelength, the unabsorbed light is seen as a mixture.
Therefore, the observed color can be explained as the complementary color of the absorbed light.
| Mainly Absorbed Light | Approximate Absorption Wavelength | Observed Color | Direction of Discussion |
|---|---|---|---|
| Violet | 400–430 nm | Yellow-green to yellow | Absorbs on the short-wavelength side |
| Blue | 430–500 nm | Orange to red | Blue light is absorbed |
| Green | 500–560 nm | Red-purple | Green light is absorbed |
| Yellow to orange | 560–620 nm | Blue | Absorbs on the long-wavelength side |
| Red | 620–750 nm | Blue-green | Red light is absorbed |
For example, if a copper(II) complex absorbs light in the orange-to-red region, the observed solution color appears blue to blue-green.
Examples of Absorption Maxima and Colors of Representative Complexes
| Complex | Central Metal | Main Ligand | Absorption Maximum λmax | Observed Color | Direction of Discussion |
|---|---|---|---|---|---|
| [Cu(H2O)6]2+ | Cu(II) | H2O | About 800 nm | Pale blue | d-d transition occurs on the long-wavelength side |
| [Cu(NH3)4]2+ | Cu(II) | NH3 | About 610 nm | Deep blue | Ligand field becomes stronger and shifts to shorter wavelength |
| [Ni(H2O)6]2+ | Ni(II) | H2O | About 395 nm, 720 nm | Green | Multiple d-d transitions |
| [Ni(NH3)6]2+ | Ni(II) | NH3 | About 360 nm, 580 nm | Blue-purple | Color changes because of ligand substitution |
| [Co(H2O)6]2+ | Co(II) | H2O | About 510 nm | Pale pink | Absorbs near the green region |
| [CoCl4]2− | Co(II) | Cl− | About 625 nm | Blue | Coordination structure and ligand change |
Even for the same metal ion, changing the ligand changes the splitting of the d orbitals, so the absorption wavelength and observed color also change.
Calculation of Absorbance Using the Beer-Lambert Law
Absorbance A is expressed using molar absorptivity ε, cell length l, and concentration c as follows.
A = εlc
| Symbol | Meaning | Unit | Example |
|---|---|---|---|
| A | Absorbance | None | 0.450 |
| ε | Molar absorptivity | L mol−1 cm−1 | 90 |
| l | Cell length | cm | 1.00 cm |
| c | Concentration | mol/L | 5.00×10−3 mol/L |
Because d-d transitions are generally not strong absorptions, their molar absorptivities are often smaller than those of charge-transfer transitions.
Example Measurements of the Relationship Between Absorbance and Concentration
The following is a reference example in which the absorbance of the [Cu(NH3)4]2+ complex was measured at 610 nm.
| Complex Concentration | Absorbance | Calculated ε | Evaluation |
|---|---|---|---|
| 1.00×10−3 mol/L | 0.090 | 90 L mol−1 cm−1 | Good |
| 2.00×10−3 mol/L | 0.181 | 90.5 L mol−1 cm−1 | Good |
| 4.00×10−3 mol/L | 0.360 | 90.0 L mol−1 cm−1 | Good |
| 6.00×10−3 mol/L | 0.535 | 89.2 L mol−1 cm−1 | Approximately linear |
| 8.00×10−3 mol/L | 0.690 | 86.3 L mol−1 cm−1 | Slightly low |
In the low-to-medium concentration range, absorbance is approximately proportional to concentration.
If linearity deviates at higher concentrations, possible causes include an excessively high concentration, incomplete complex formation, scattering, or contamination of the cell.
Example Calculation of Molar Absorptivity
If the complex concentration is 5.00×10−3 mol/L, the cell length is 1.00 cm, and the absorbance is 0.450,
ε = A / lc
ε = 0.450 ÷ (1.00 × 5.00×10−3) = 90 L mol−1 cm−1
This value corresponds to a relatively weak absorption for a d-d transition.
Even for the same complex, if the measurement wavelength is shifted away from the absorption maximum, the apparent molar absorptivity becomes smaller.
Changes in Absorption Maxima Caused by Ligands
When the ligand changes, the splitting of the d orbitals of the metal ion changes.
When the splitting becomes larger, higher-energy light is absorbed, so the absorption wavelength becomes shorter.
| Cu(II) Complex | Main Ligand | Absorption Maximum | Observed Color | Direction of Discussion |
|---|---|---|---|---|
| [Cu(H2O)6]2+ | H2O | About 800 nm | Pale blue | Relatively weak ligand field |
| [Cu(NH3)4]2+ | NH3 | About 610 nm | Deep blue | Ligand field becomes stronger |
| Cu(II)-en complex | Ethylenediamine | About 580 nm | Blue-purple | Further toward shorter wavelength |
When the ligand changes from water to ammonia or ethylenediamine, there are examples in which the absorption maximum shifts toward shorter wavelengths.
This can be explained by an increase in ligand-field splitting.
Ligand-Field Strength and Absorption Wavelength
| Ligand Field | d-Orbital Splitting | Energy of Absorbed Light | Absorption Wavelength | Direction of Discussion |
|---|---|---|---|---|
| Weak | Small | Low | Long-wavelength side | More likely to absorb red to near-infrared light |
| Moderate | Moderate | Moderate | Near the center of the visible region | Color changes are readily observed |
| Strong | Large | High | Short-wavelength side | More likely to absorb blue to violet light |
The energy of light is inversely proportional to wavelength.
Therefore, if the absorption maximum shifts toward shorter wavelengths, the splitting of the d orbitals is considered to have increased.
Example of Determining Transition Energy From Absorption Wavelength
The energy corresponding to a d-d transition can be approximately calculated from the absorption maximum.
E = hc / λ
When light of λ = 610 nm is absorbed, the energy per mole can be approximately calculated as follows.
E = NAhc / λ ≒ 196 kJ/mol
| Absorption Maximum | Region of Light | Transition Energy | Direction of Discussion |
|---|---|---|---|
| 800 nm | Red to near-infrared | About 150 kJ/mol | Small splitting |
| 610 nm | Orange to red | About 196 kJ/mol | Moderate splitting |
| 580 nm | Yellow to orange | About 206 kJ/mol | Slightly larger splitting |
| 450 nm | Blue to violet | About 266 kJ/mol | Large splitting |
The shorter the absorption wavelength, the greater the energy required for the transition.
Example of Changes in the Absorption Spectrum Caused by Complex Formation
When a ligand is added to a metal ion, the type of complex changes, and the absorption spectrum and solution color also change.
| Condition | Main Complex Species | Absorption Maximum | Absorbance | Observed Color | Direction of Discussion |
|---|---|---|---|---|---|
| Cu2+ aqueous solution | Aqua complex | About 800 nm | 0.180 | Pale blue | Water is the main ligand |
| Small amount of NH3 added | Mixed-ligand complex | About 700 nm | 0.300 | Blue becomes deeper | Ligand substitution proceeds |
| Excess NH3 added | Ammine complex | About 610 nm | 0.520 | Deep blue | Complex formation proceeds |
| Precipitation caused by pH change | Hydroxide or similar species | Unclear | Appears larger because of scattering | Turbid | Unsuitable for absorbance measurement |
If the absorption maximum shifts and the absorbance changes when a ligand is added, the composition or coordination environment of the complex is considered to have changed.
Examples When Absorbance Is Not Proportional to Concentration
| Cause | Observed Phenomenon | Effect on Absorbance | Direction of Discussion |
|---|---|---|---|
| Incomplete complex formation | Insufficient ligand | Absorbance lower than expected | Not all metal ions have formed the same complex |
| Precipitation or turbidity | Solution becomes opaque | Appears high because of scattering | Effect of scattering rather than absorption |
| Concentration too high | Absorbance exceeds 1 | Linearity becomes poor | Dilution is necessary |
| Insufficient blank correction | Absorption of solvent or ligand remains | Absorbance is overestimated | Use an appropriate blank |
| Dirty cell or bubbles | Measured values fluctuate | Absorbance becomes unstable | Clean the cell and remove bubbles |
Example Organization of Complex Color and Absorption Spectra
| Observed Color | Main Absorption Wavelength | Absorbed Color | Color Seen as Complementary Color | Direction of Discussion |
|---|---|---|---|---|
| Blue | 600–700 nm | Orange to red | Blue | Red light is absorbed |
| Green | Near violet and red | Multiple wavelengths are absorbed | Green remains | Multiple d-d transitions |
| Red-purple | 500–560 nm | Green | Red-purple | Green light is absorbed |
| Yellow | 400–450 nm | Violet to blue | Yellow | Absorbs on the short-wavelength side |
The actual color of a complex is affected not only by a single absorption but also by multiple absorption bands, concentration, optical path length, and background.
Example of Variation in Measured Values
The following is an example of variation in the absorption maximum and absorbance when the same complex solution is measured multiple times.
| Measurement | λmax | Absorbance | Evaluation | Possible Factor |
|---|---|---|---|---|
| 1st | 610 nm | 0.520 | Good | Reference |
| 2nd | 611 nm | 0.518 | Good | Within measurement error |
| 3rd | 610 nm | 0.526 | Good | Slight difference in cell position |
| 4th | 625 nm | 0.610 | Possible outlier | Bubble, dirty cell, turbidity |
If only the fourth measurement shows a large deviation in both the absorption maximum and absorbance, contamination of the cell, bubbles, or scattering caused by precipitation should be suspected rather than an intrinsic change in the complex.
Example of How to Write the Results
When ammonia was added to an aqueous Cu(II) solution, the solution color changed from pale blue to deep blue.
Measurement of the absorption spectrum showed absorption around 800 nm for the aqua complex, whereas an absorption maximum around 610 nm was observed for the ammine complex.
From this, the coordination environment of Cu(II) is considered to have changed because the ligand changed from water to ammonia.
The shift of the absorption maximum from 800 nm to 610 nm toward the shorter-wavelength side indicates that the energy required for the d-d transition increased.
Because the energy of light is inversely proportional to wavelength, the shift toward shorter wavelengths corresponds to an increase in the splitting of the d orbitals.
Therefore, the ammonia ligand is considered to produce a stronger ligand field than the water ligand.
The observed color of a complex can be explained as the complementary color of the absorbed light.
Because the [Cu(NH3)4]2+ complex absorbs light in the orange-to-red region, it is considered to appear deep blue as a combination of the remaining light.
In this way, the cause of the color of a complex can be explained by relating the position of the absorption-spectrum peak to the observed color.
Points for Connecting the Results to the Discussion
In discussing the absorption spectrum of a complex, it is important to relate the position of the absorption maximum, observed color, ligand-field strength, complex formation, and measurement errors.
- Can you read the absorption maximum λmax and explain which wavelength of light is being absorbed?
- Can you explain the observed color as the complementary color of the absorbed light?
- Can you explain that d-d transitions originate from splitting of the d orbitals caused by the ligand field?
- Do you understand that shorter absorption wavelengths correspond to greater transition energies?
- Can you discuss how changing the ligand changes the d-orbital splitting and therefore changes the color and absorption spectrum?
- Can you explain the relationship among absorbance, concentration, and molar absorptivity using the Beer-Lambert law?
- If absorbance is not proportional to concentration, can you consider incomplete complex formation, precipitation, scattering, and insufficient blank correction?
- Can you explain the effects of cell contamination, bubbles, pH changes, and turbidity on the absorption spectrum?
Example Discussion
In this experiment, the absorption spectra of Cu(II) complexes were measured to investigate changes in color and absorption maxima caused by different ligands.
The aqueous Cu(II) solution was pale blue and showed absorption around 800 nm.
In contrast, after ammonia was added, the solution became deep blue and the absorption maximum shifted to around 610 nm.
From this change, the coordination environment of Cu(II) is considered to have changed because of ligand substitution.
In transition-metal complexes, the d orbitals of the metal ion are split by the electric field of the ligands.
When a d electron moves from a lower-energy d orbital to a higher-energy d orbital, light of the corresponding wavelength is absorbed.
In this experiment, the shift of the absorption maximum from 800 nm to 610 nm toward the shorter-wavelength side indicates that the energy required for the transition increased.
Therefore, the splitting of the d orbitals is considered to have been larger with the ammonia ligand than with the water ligand.
The observed color can be explained as the complementary color of the absorbed light.
The ammine copper(II) complex appears deep blue because it absorbs light in the orange-to-red region, leaving the blue component relatively strong.
However, the actual color is affected not only by a single absorption band but also by multiple absorption bands, concentration, cell length, and observation conditions.
When the relationship between absorbance and concentration was examined, absorbance was approximately proportional to concentration in the low-to-medium concentration range according to the Beer-Lambert law.
However, at higher concentrations, a slight deviation from linearity was observed.
Possible causes include excessively high absorbance, incomplete complex formation, and scattering caused by turbidity of the solution or contamination of the cell.
Possible sources of measurement error include insufficient blank correction, contamination or bubbles on the cell surface, errors in concentration preparation, changes in complex species caused by pH changes, and formation of precipitates.
In particular, when precipitation or turbidity is present, the measured absorbance is affected not only by light absorption but also by scattering, so the spectrum shape and absorbance become inaccurate.
Therefore, when discussing the color and absorption maximum of a complex, it is necessary to confirm that the solution is uniform and transparent.
Summary
In the absorption spectrum of a complex, the energy of the d-d transition and the strength of the ligand field can be discussed from the wavelength of the absorption maximum.
When the ligand changes, the splitting of the d orbitals changes, and the absorption wavelength and observed color also change.
This reference example covered absorption maxima of representative complexes, complex color and complementary colors, the Beer-Lambert law, molar absorptivity, changes in absorption wavelength caused by ligands, transition energy, spectral changes caused by complex formation, and measurement errors.
In a report, it is useful to discuss the absorption maximum, observed color, ligand field, and measurement conditions in relation to one another.
What Is a d-d Transition?
A d-d transition is an electronic transition in which an electron in a d orbital of a transition-metal ion moves to another d orbital split by the ligand field.
In transition-metal complexes, when ligands bind around the metal ion, the energies of the d orbitals are no longer equal.
As a result, lower-energy and higher-energy d orbitals are produced.
When light corresponding to this energy difference is absorbed, an electron transitions from a lower d orbital to a higher d orbital.
If this absorption lies in the visible region, the complex shows color.
Example Discussion:
An absorption peak was observed in the visible region for the measured complex.
This is considered to have occurred because the d orbitals of the central metal ion were split by the ligand field and light corresponding to the energy difference was absorbed, causing a d-d transition.
Therefore, the position of the absorption peak is considered to reflect the magnitude of the ligand field in the complex.
Relationship Between Ligand-Field Splitting and Absorption Wavelength
Ligand-field splitting refers to the splitting of the energies of the d orbitals of a metal ion caused by the influence of ligands.
The larger the splitting, the greater the energy required for an electronic transition.
Because light energy increases as wavelength becomes shorter, complexes with larger ligand-field splitting may show absorption shifted toward shorter wavelengths.
In contrast, if the ligand-field splitting is small, less energy is required for the electronic transition and absorption is more likely to appear at longer wavelengths.
Therefore, the wavelength of the absorption peak provides an important clue for considering ligand-field strength.
| Ligand-Field Splitting | Required Light Energy | Trend in Absorption Wavelength |
|---|---|---|
| Large | Large | Tends toward shorter wavelengths |
| Small | Small | Tends toward longer wavelengths |
Example Discussion:
Changing the ligand caused the position of the absorption peak to change.
This is considered to be because the splitting of the d orbitals of the metal ion changed depending on the type of ligand.
When ligand-field splitting is large, more energy is required for the d-d transition and the absorption tends to shift toward shorter wavelengths.
Therefore, the change in absorption wavelength is considered to reflect the difference in ligand-field strength.
Relationship Between the Color of a Complex and the Light It Absorbs
The color of a complex is not the color of the light absorbed by the complex itself, but the color of the light that is transmitted or reflected without being absorbed.
For example, if a complex strongly absorbs light near the red region, the observed color may be close to its complementary color.
In other words, by checking where peaks occur in the visible region of the absorption spectrum, it is possible to explain why a complex appears to have a particular color.
| Mainly Absorbed Light | Example of Visible Color |
|---|---|
| Near violet to blue | Yellow to orange |
| Near blue to green | Red to purple |
| Near green to yellow | Red-purple to blue-purple |
| Near orange to red | Blue to green |
Example Discussion:
The complex solution appeared blue not because it absorbed blue light, but because mainly blue light was transmitted or reflected and observed.
If the absorption spectrum shows absorption in the orange-to-red region, the complex may appear blue, which is close to the complementary color.
Therefore, the observed color of the complex corresponds to the absorption wavelength in the visible region of the absorption spectrum.
Discussion of the Maximum Absorption Wavelength
The maximum absorption wavelength is the wavelength at which absorbance is greatest.
It is also expressed as λmax.
The maximum absorption wavelength of a complex changes depending on the central metal, oxidation state, ligand, and coordination structure.
In a report, discuss where the maximum absorption wavelength lies, whether it is close to the literature value or standard sample, and in which direction it shifts when the ligand is changed.
Example Discussion:
The maximum absorption wavelength of the measured complex was observed around ○○ nm.
This wavelength is considered to correspond to absorption of light matching the energy difference between the d orbitals of the central metal.
If the maximum absorption wavelength is close to the literature value or that of a standard sample, this supports the possibility that the target complex was formed.
On the other hand, if it differs greatly, the ligand environment or oxidation state may differ from that of the intended complex.
What Can Be Learned From the Magnitude of Absorbance?
Absorbance is a value indicating how much light is absorbed by a sample.
For the same complex, higher concentrations tend to give higher absorbance.
In addition, the longer the cell path length, the greater the distance traveled by light through the sample and therefore the larger the absorbance.
A = εcl
Here, A is absorbance, ε is molar absorptivity, c is concentration, and l is the cell length.
This relationship is also used when determining concentration by spectrophotometry.
Example Discussion:
When compared at the same wavelength, samples with higher concentrations showed greater absorbance.
This is because more complex molecules or complex ions that absorb light were present in the solution, causing more of the transmitted light to be absorbed.
Therefore, the magnitude of absorbance is considered to reflect differences in complex concentration.
The Beer-Lambert Law and Complex Concentration
When absorption spectra of complexes are treated quantitatively, the relationship between absorbance and concentration can be explained using the Beer-Lambert law.
When measurements are performed at the same wavelength and with the same cell length, there is a range in which absorbance is proportional to complex concentration.
In experiments in which a calibration curve is prepared, this proportional relationship is used to determine the concentration of an unknown sample.
Example Discussion:
If absorbance increased linearly with concentration, the Beer-Lambert law can be considered to hold approximately within the measurement range.
On the other hand, if the higher-concentration values deviated from the straight line, possible causes include excessively high absorbance reducing measurement accuracy, interactions between complexes, stray light, and errors in concentration preparation.
Discussion When a Peak Is Broad
Absorption peaks of complexes may be sharp or broad.
Absorption caused by d-d transitions may appear broad because of molecular vibrations, interactions with the solvent, or the presence of multiple coordination environments.
In addition, if multiple absorptions overlap, the peak may broaden or a shoulder peak may appear.
Example Discussion:
One possible reason the observed absorption peak was broad is that the d-d transition was affected by vibrational states and the solvent environment.
In addition, if complexes with multiple coordination environments were present in the solution, absorptions at similar wavelengths may have overlapped and caused the peak to appear broad.
Therefore, the peak shape provides a clue for discussing the uniformity of the complex and differences in coordination environment.
Discussion When Multiple Peaks Are Present
If multiple peaks appear in the absorption spectrum, multiple electronic transitions may be occurring.
Possible cases include multiple d-d transitions, simultaneous observation of d-d and charge-transfer transitions, or coexistence of different complexes.
Example Discussion:
The presence of multiple peaks in the absorption spectrum suggests that multiple electronic transitions may be occurring in the complex.
One peak may originate from a d-d transition, while another may originate from a charge-transfer transition from ligand to metal or from metal to ligand.
The coexistence of multiple complex species in the sample may also cause multiple absorption peaks to be observed.
Difference From Charge-Transfer Transitions
Absorption by complexes may involve not only d-d transitions but also charge-transfer transitions.
A charge-transfer transition is an electronic transition in which electron density moves from ligand to metal or from metal to ligand.
In general, charge-transfer transitions may be observed as stronger absorptions than d-d transitions.
Therefore, if a very strong absorption peak is observed, it is useful to consider not only a d-d transition but also the possibility of a charge-transfer transition.
| Type of Transition | Characteristics | Perspective for Discussion |
|---|---|---|
| d-d transition | Electronic transition between d orbitals of the metal ion | Related to ligand-field splitting and the color of the complex |
| Charge-transfer transition | Transition involving electron transfer between the metal and ligand | Tends to appear as strong absorption |
Example Discussion:
If the observed absorption was very strong, the contribution of a charge-transfer transition in addition to a d-d transition may be considered.
A d-d transition originates from d-orbital splitting caused by the ligand field, whereas a charge-transfer transition involves movement of electron density between the metal and ligand.
Therefore, the type of transition must be considered carefully for peaks with large absorption intensity.
Spectral Changes When the Ligand Is Changed
Changing the ligand changes the absorption spectrum of the complex.
This is because the strength of the ligand field and the interaction with the metal ion change.
When the ligand field changes, the energy splitting of the d orbitals changes, and the absorption wavelength corresponding to the d-d transition also changes.
Therefore, even with the same metal ion, the positions of absorption peaks and the color of the complex may change when the ligand is water, ammonia, ethylenediamine, chloride ion, or another species.
Example Discussion:
Changing the ligand caused the maximum absorption wavelength to change.
This is considered to be because the ligand-field splitting of the metal ion changed depending on the type of ligand.
When the ligand field becomes stronger, the energy required for the d-d transition increases and the absorption may shift toward shorter wavelengths.
This change in absorption wavelength is also reflected in the difference in the color of the complex.
Spectral Changes When the Oxidation State Is Changed
When the oxidation state of the central metal changes, the number of d electrons and the interaction between the metal and ligands change.
As a result, the position and intensity of absorption peaks and the color of the complex may change.
In iron and cobalt complexes, for example, the spectrum may change because of redox reactions.
Example Discussion:
Because the absorption spectrum changed after the redox treatment, the oxidation state of the central metal may have changed.
When the oxidation state changes, the number of d electrons and the condition of the ligand field change, and the energy required for electronic transitions also changes.
Therefore, changes in the position and intensity of absorption peaks provide a clue indicating changes in the oxidation state of the metal ion.
Relationship Between Coordination Structure and Absorption Spectrum
The coordination structure of a complex also affects its absorption spectrum.
The manner of d-orbital splitting differs among octahedral, tetrahedral, square-planar, and other structures.
Therefore, even with the same metal ion and oxidation state, changing the structure may change the absorption wavelength or number of peaks.
In a report, the structure should not be completely determined from the absorption spectrum alone, but should be estimated together with color, magnetic properties, composition, literature values, and similar information.
Example Discussion:
When the coordination structure of a complex changes, the way in which the d orbitals split also changes.
Therefore, octahedral and tetrahedral complexes may show different absorption spectra even when they contain the same metal ion.
The absorption peaks in this experiment are not inconsistent with the expected coordination structure of the target complex, but additional information such as magnetic properties and infrared absorption spectra is also necessary to determine the structure completely.
Example Discussion of the Absorption Spectrum of a Copper Complex
Copper(II) complexes may show relatively broad absorption in the visible region.
When the coordination environment of the copper(II) ion changes from water to ammonia, ethylenediamine, or another ligand, the color and absorption peak of the complex change.
Blue copper complexes may absorb light in the red-to-orange region.
Example Discussion:
In the absorption spectrum of the copper complex, broad absorption was observed in the visible region.
This is considered to have occurred because the d orbitals of the copper(II) ion were split by the ligand field and a d-d transition occurred.
Because the solution appeared blue, light in the orange-to-red region of the visible spectrum is considered to have been absorbed, and a blue color close to the complementary color was observed.
Example Discussion of the Absorption Spectrum of a Nickel Complex
In nickel(II) complexes, the absorption spectrum and color change greatly depending on the coordination structure and ligands.
Differences among octahedral, tetrahedral, square-planar, and other structures also affect d-orbital splitting and magnetic properties.
Therefore, in nickel complexes, it is easy to discuss absorption spectra in relation to color and magnetic properties.
Example Discussion:
The absorption peaks of the nickel complex are considered to reflect differences in d-orbital splitting caused by the type of ligand and coordination structure.
Nickel(II) has a d8 electron configuration, so the electronic state and magnetic properties change depending on the coordination structure.
By considering the changes in the absorption spectrum together with the magnetic results, the coordination structure of the complex can be discussed more specifically.
Example Discussion of the Absorption Spectrum of an Iron Complex
In iron complexes, the absorption spectrum changes greatly depending on the oxidation state of iron and the type of ligand.
Because Fe(II) and Fe(III) differ in the number of d electrons and in their interactions with ligands, their absorption wavelengths and colors may differ even with the same ligand.
In color reactions such as thiocyanate complexes, absorbance measurements may also be used to evaluate iron concentration.
Example Discussion:
Because the absorption spectrum of the iron complex changed as a result of redox treatment, the oxidation state of the iron may have changed.
Fe(II) and Fe(III) have different numbers of d electrons, and their interactions with ligands and the energies of electronic transitions also differ.
Therefore, the change in absorption peaks is considered to reflect changes in the oxidation state and coordination environment of the iron complex.
Example Discussion of the Absorption Spectrum of a Cobalt Complex
In cobalt complexes, the color changes greatly depending on differences between the Co(II) and Co(III) oxidation states and the type of ligand.
In addition, differences in coordination structure, such as octahedral and tetrahedral forms, may change the energy and absorption intensity of d-d transitions.
Example Discussion:
The peaks observed in the absorption spectrum of the cobalt complex may originate from d-d transitions of the central metal.
When the ligand or oxidation state changes, the d-orbital splitting and electron configuration of the cobalt ion change, and the absorption wavelength also changes.
Therefore, differences in the spectrum provide a clue to differences in the coordination environment and oxidation state of the complex.
Errors When the Concentration Is Too High
If the sample concentration is too high, the absorbance becomes too large and accurate measurement may become difficult.
In a region where absorbance is too high, transmitted light becomes weak and the influence of instrumental errors and stray light becomes larger.
In addition, at high concentrations, interactions or association between complexes may occur and the absorption spectrum may change.
Example Discussion:
If the absorbance was too large, the transmitted light may have become weak and the effects of instrumental readout errors and stray light may have increased.
In addition, at high concentrations, interactions between complexes may change the positions and shapes of absorption peaks.
Therefore, to measure the absorption spectrum accurately, it is important to dilute the sample to an appropriate concentration.
Errors When the Concentration Is Too Low
If the sample concentration is too low, the absorbance becomes small and peaks are more easily buried in noise.
As a result, it becomes difficult to accurately read the maximum absorption wavelength and absorbance.
Especially for complexes with weak d-d absorption, peaks may become unclear when the concentration is too low.
Example Discussion:
One possible reason the absorption peak was unclear is that the sample concentration was too low.
At low concentrations, absorbance becomes small and the peak is more easily buried in instrumental noise or baseline fluctuations.
Therefore, to accurately evaluate the peak position and absorbance, the sample must be adjusted to a concentration range suitable for measurement.
Importance of Blank Correction
In absorption-spectrum measurements, a blank measurement is performed to remove absorption caused by the solvent or the cell itself.
If blank correction is insufficient, absorption caused by the solvent or cell rather than the sample is included in the spectrum.
As a result, the baseline may shift or the absorbance may appear larger than it actually is.
Example Discussion:
One possible reason the baseline was shifted is that blank correction was insufficient.
Because the solvent and cell themselves also have slight effects from absorption and reflection, these effects must be corrected using a blank.
If blank correction is inappropriate, absorption not originating from the sample complex is included, causing errors in the evaluation of absorbance and peak shape.
Errors Caused by Cell Contamination or Bubbles
In spectroscopic measurements, contamination, fingerprints, scratches, bubbles, and differences in liquid level in the cell affect measured values.
Contamination on the cell surface and bubbles scatter light and can cause the absorbance to appear larger than it actually is.
Example Discussion:
One possible reason the absorbance was larger than expected is scattering of light caused by contamination on the cell surface or bubbles.
If fingerprints or contamination are present on the cell, scattering and reflection occur in addition to absorption by the sample.
Bubbles in the solution also disturb the optical path, so it is important to clean the cell and remove bubbles before measurement.
Effect of the Solvent
The solvent may affect the coordination environment and absorption spectrum of a complex.
If solvent molecules coordinate to the metal ion, they may compete with the target ligand and change the structure or color of the complex.
In addition, the position and shape of absorption peaks may change depending on solvent polarity and hydrogen-bonding properties.
Example Discussion:
Changing the measurement solvent may change the position and shape of the absorption peaks.
If solvent molecules coordinate to the metal ion, the coordination environment of the complex changes and may also affect the energy splitting of the d orbitals.
In addition, differences in solvent polarity and solvation may change the energy of electronic transitions, so solvent conditions must be kept the same when comparing spectra.
When the Complex Changes During Measurement
A complex may change during measurement because of the passage of time, light, redox reactions, ligand exchange, or pH changes.
If the color of the solution changes before and after measurement or if the absorbance changes over time, the stability of the complex must be considered.
Example Discussion:
If the absorption spectrum changed during measurement, the coordination environment or oxidation state of the complex may have changed over time.
If ligand exchange or a redox reaction proceeds, complex species with different absorptions are produced and the peak positions and absorbance change.
Therefore, measurements should be performed under conditions kept as consistent as possible after solution preparation, and if time-dependent changes occur, their effects must be considered.
Discussion When the Results Differ From Literature Values
If the measured maximum absorption wavelength or absorbance differs from the literature value, possible causes include concentration, solvent, pH, temperature, incomplete ligand substitution, impurities, and instrumental conditions.
When comparing with literature values, it is important to confirm whether the measurement conditions are the same.
Example Discussion:
One possible reason the measured maximum absorption wavelength differed from the literature value is that the measurement solvent or pH conditions were different.
The absorption spectrum of a complex is sensitive to its coordination environment, and changes in solvent or pH may alter the ligand-binding state or complex species.
In addition, the presence of impurities or unreacted substances other than the target complex may also affect the peak position and shape.
When the Result Can Be Considered Good
A good result in absorption-spectrum measurement of a complex is indicated when the peaks are clear, the baseline is stable, the maximum absorption wavelength does not greatly contradict the literature value or standard sample, and the relationship between the observed color and absorption wavelength can be explained.
In addition, if the relationship between concentration and absorbance is linear, the reliability of the measurement as a quantitative analysis is also increased.
Example Discussion:
In the measured absorption spectrum, a clear absorption peak was observed in the visible region, and the maximum absorption wavelength generally agreed with the literature value for the target complex.
In addition, the observed color of the sample solution could be explained as the complementary color of the absorbed light.
From this, the measurement result is considered to reflect absorption originating from a d-d transition of the target complex.
Example Discussion When the Measurement Did Not Go Well
When absorption-spectrum measurement does not go well, possible causes can be considered from results such as unclear peaks, absorbance that is too high or too low, a shifted baseline, disagreement with literature values, or inconsistency between color and spectrum.
It is easier to organize the discussion by separately considering concentration, blank correction, the condition of the cell, solvent, pH, complex stability, and impurities.
Example Discussion:
One possible reason the absorption peak was unclear is that the sample concentration was too low.
When the concentration is low, absorbance becomes small and the peak is more easily buried in noise or baseline fluctuations.
In addition, if blank correction was insufficient or bubbles or contamination were present in the cell, the spectrum shape may have been disturbed and the peak position may have become difficult to determine.
How to Write Points for Improvement
In a discussion of the absorption spectrum of a complex, including points for improvement as well as sources of error makes the report easier to organize.
Improvements are easier to write when divided into sample preparation, measurement operations, and data processing.
Improvements to Sample Preparation
- Dilute to an appropriate concentration
- Measure under pH conditions where the complex is stable
- Mix sufficiently before measurement
- Remove insoluble matter and precipitates before measurement
- Keep the time until measurement consistent
- Use solvent conditions consistent with standard samples or literature values
Improvements to Measurement Operations
- Perform blank measurement correctly
- Wipe contamination and fingerprints from the cell
- Remove bubbles before measurement
- Keep the orientation of the cell consistent
- Dilute the sample if the absorbance is too high
- Compare samples over the same wavelength range and under the same conditions
Improvements to Data Processing
- Read the maximum absorption wavelength accurately
- Check for baseline shifts
- Consider the possibility of overlapping peaks
- When comparing with literature values, also check the measurement conditions
- Relate color, ligand, structure, and spectrum
Example of How to Write Points for Improvement:
To measure the absorption spectrum more accurately, the sample concentration must be adjusted to an appropriate range so that the absorbance is neither too high nor too low.
In addition, correctly performing blank correction and removing contamination or bubbles from the cell can reduce the effects of absorption and scattering not caused by the sample.
Furthermore, if the complex may change over time, it is important to keep the time from solution preparation to measurement consistent.
Difference Between a Superficial Discussion and a Good Discussion
In a discussion of the absorption spectrum of a complex, simply writing that “there was a peak” or “it was blue” results in a superficial discussion.
A persuasive discussion can be produced by connecting d-d transitions, ligand-field splitting, absorption wavelength, complementary colors, and differences in ligands and structures.
| Superficial Discussion | Good Discussion |
|---|---|
| A peak appeared. | Because an absorption peak was observed in the visible region, the d orbitals of the central metal ion are considered to have been split by the ligand field, and light corresponding to the energy difference was absorbed, causing a d-d transition. |
| It was blue. | The complex appeared blue because mainly blue light was transmitted or reflected and observed. If the absorption spectrum shows absorption in the orange-to-red region, the complex is considered to appear blue, which is close to the complementary color. |
| It differed from the literature value. | Possible reasons the maximum absorption wavelength differed from the literature value include differences in solvent, pH, ligand environment, oxidation state, impurities, and concentration conditions. Because the absorption spectrum of a complex is sensitive to the coordination environment, measurement conditions must be kept consistent when making comparisons. |
Examples of Expressions That Can Be Used in Reports
The following expressions can be used when writing the results and discussion of absorption spectra of complexes.
Adjust the necessary parts according to your own experimental results.
- Because an absorption peak was observed in the visible region, an electronic transition is considered to have occurred in the complex.
- This absorption may originate from a d-d transition occurring between d orbitals of the central metal ion.
- The d orbitals are considered to have been split by the ligand field, and light corresponding to the energy difference was absorbed.
- Because the complementary color of the absorbed light is observed as the visible color, the absorption wavelength corresponds to the color of the complex.
- When the type of ligand changes, the magnitude of ligand-field splitting changes and the maximum absorption wavelength also changes.
- When the oxidation state changes, the number of d electrons and coordination environment change, affecting the absorption spectrum.
- If the absorption is very strong, a charge-transfer transition in addition to a d-d transition may be involved.
- If the absorbance is too high, the effects of stray light and measurement error may become large.
- If blank correction is insufficient, absorption caused by the solvent or cell may be included in the measured result.
- Because the structure cannot be completely determined from color and absorption spectra alone, it must be judged together with results such as magnetic properties and infrared absorption spectra.
Points to Check When Discussing the Absorption Spectrum of a Complex
Checking the following points before writing the report makes the discussion easier to write.
- Have you recorded the maximum absorption wavelength?
- Have you recorded the magnitude of the absorbance?
- Have you related the color of the sample solution to the absorption wavelength?
- Have you explained the meaning of a d-d transition?
- Have you discussed the relationship between ligand-field splitting and absorption wavelength?
- Have you explained spectral changes caused by differences in ligands?
- Have you considered the effects of oxidation state and coordination structure?
- Have you considered the possibility of charge-transfer transitions?
- Have you checked whether the concentration was appropriate?
- Have you considered blank correction and cell contamination as sources of error?
- When comparing with literature values, have you considered differences in measurement conditions?
- Have you avoided determining the structure from the spectrum alone?
Summary
In the absorption spectrum of a complex, the wavelengths of light absorbed by the complex are measured and used to discuss color, electronic state, and coordination environment.
In transition-metal complexes, d-d transitions in which electrons move between d orbitals split by the ligand field may be observed as absorption in the visible region.
The color of a complex is observed not as the color of the absorbed light itself, but as the light that remains unabsorbed or as its complementary color.
Therefore, it is important to explain the relationship between the positions of peaks in the absorption spectrum and the color of the sample solution.
When the type of ligand, oxidation state of the central metal, or coordination structure changes, the splitting of the d orbitals changes, and the maximum absorption wavelength and color also change.
In a report, do not simply write that “a peak appeared” or “the color changed.”
Discuss the results by relating d-d transitions, ligand-field splitting, absorption wavelength, complementary colors, and differences in ligands and oxidation states.
In addition, considering sources of error such as concentration, blank correction, contamination of the cell, solvent, pH, and impurities produces a persuasive discussion of the absorption spectrum.
