UV-Vis spectroscopy is one of the most important quantitative spectroscopic techniques.  The wavelength range extends from about 190 nm to 750 nm which corresponds to electronic transitions of different origins.  The energy level diagram can simply be represented as shown in Figure 1.




      Figure 1:  Energy level diagram representing electronic transitions.


      Four transitions are shown in the figure, which correspond to:


      a.  s - s*  Transitions


           This type of transition requires large energy which may result in breakdown of chemical bonds.  It is not important from the analytical point of view for reasons to be discussed later.  The radiation wavelength which should be used to enforce this type of transition is below 190 nm .


      b. p - p* Transitions


           This is the most analytically useful type of transitions where radiation of certain wavelength cause electrons present in the p bonds to be excited to the higher energy p* state.  The energy required is moderate and can be obtained from sources operating in the UV-Vis range.


      c. n - p*  and  n - s*  Transitions


           Molecules which contain lone pair(s) of electrons exhibit some special characteristics.  Electrons which do not participate in chemical bonds can absorb energy and are excited either to the p* (if the molecule has p bonds) or to the   s*  state.  Energy required for a n - p* transition is small while a n - s*  transition requires more energy.  Factors to be mentioned later eliminate the use of these transitions in analytical work.


Beer's Law


           It is logical that the amount of energy absorbed by molecules depends on the number of the absorbing species and, therefore, on concentration.  A quantitative relationship was derived by Beer, Lambert and Bouger separately.  Therefore, the relationship can be referred to as Beer's, lambert or Bouger law or some combination of these names.


      The law states that:


      A =  ebC, where


      A  is the absorbance.

      b  is the cell length in cm

      e  is the molar absorptivity in L cm-1 mol-1

      C  is the concentration in mol/L

      Away from mathematical derivation, the law can be simply derived as follows.


      It is logical to assume that A a C since as C increases, the number of absorbing species increases and thus A increases.

      Also A a b since as b increases more molecules are encountered by radiation and therefore possibility of absorbance increases.


      This can yield the relation

      A   a  bC                  or

      A = ebC, where e is a proportionality constant and is called molar absorptivity.  Old conventions used the expression “extinction coefficient” or “optical density” for the molar absorptivity.  However, the last expression is the internationally accepted one.


      In Beer's law, e is the most important sensitivity indicator.  When b is 1 cm and A is plotted as a function of the concentration, a straight line relationship is obtained as in Figure 2.


      Figure 2:  A straight line relationship between absorbance (A) and concentration (C). Molar absorptivity is the slope when b equals 1 cm.


           When the concentration is expressed in mol/L and is plotted on the X-axis, e is the slope and its value is an indication of the sensitivity of the method.


Deviations from Beer's Law


           Beer's law suggests direct proportionality between Absorbance and concentration and that a straight line relationship should be obtained.  However, some factors can lead to different behavior.


      1.  Dilute Solutions


           It is observed that only solutions that are less than 0.01 M can result is successful application of Beer's law.  High concentrations of solute lead to interactions between neighboring molecules resulting in a change in the absorbance characteristics of molecules.


      2.  Monochromatic Light


           As the energy between any two energy levels is constant, one can roughly say that the wavelength of the radiation that can cause an electronic transition from one energy level to the other should be exact.  Therefore, Beer's law is valid only when a monochromatic light is used. This can be validated through a simple mathematical treatment where only when the molar absorptivities of the two wavelengths are equal that beer's law holds, a condition which will never be attained.


      3.  Stray Radiation


           In some cases, instrumental artifacts can lead to considerable amount of stray radiation (Radiation reaching the detector without passing through the sample).


                      Po  +  Ps

      A =  log   _______

                        P   +   Ps


           At high concentrations, P is small and absorbance is affected strongly by the value of Ps resulting in a nonlinear relationship between absorbance and concentration.


      4.  Refractive Index


           The application of Beer's law requires the measurement of the absorbance of different concentrations.  Since  e  is dependent on refractive index, it is expected that solutions for which the refractive index changes a lot with concentration will show a deviation from Beer's law.


      5.  Chemical Reactions or Transformations


           If an analyte reacts with any species or dissociates forming species of different absorption characteristics, the absorbance may not necessarily be as predicted by Beer's law.


The Molar Absorptivity


           From the above argument, it is clear that e  is an important element in  Beer's law expression.  It is an indicator of sensitivity which means that the value of e   is very important in characterizing a system for quantitative capacity.  As  e   increases, it becomes easier to determine lower concentrations of analytes.

           The molar absorptivity has large values for p - p*  transitions ranging from 1000 to 10000 L cm-1 mol-1.   For n - p* transitions, e ranges from 10 - 100 L cm-1 mol-1.  This means that the most important trasition in UV-Vis is the p  - p* transition and, therefore, will be subjected to further studies in different solvents.


Effects of Solvent Polarity on p  - p*  Transition


           Solvent polarity is an important factor in the definition of the energy required to cause a p  - p* transition.  Figure 3 shows the effect of polar solvents on the energy of this transition



           Nonpolar solvent                         Polar solvent


      Figure  3 :  Effect of Solvent Polarity on  p  - p* Transitions.


      p* is more polar than p and is stabilized more than p in polar solvents.  The energy required for p  - p* transition in polar solvents is thus reduced and the wavelength of incident radiation increases. This is referred to as bathochromic shift or red shift.


      A n - p* transition is affected in an opposite way since the n electrons are stabilized more than the p* leading to increased energy and shorter wavelength. This is referred to as hypsochromic shift or blue shift.


Qualitative Versus Quantitative Analysis


           UV-Vis is mainly used in quantitative analysis due to the fact that spectra obtained from this type of spectroscopic technique are almost featureless.  It is very difficult to assign a spectrum for a specific species since a spectrum has no definite characteristics associated with a species.  However, Beer's law can be efficiently applied for quantitative determination of any species which absorbs in the UV-Vis region.  Therefore, UV-Vis spectroscopy is mainly a quantitative technique and is marginally used in qualitative analysis.


Additivity of Absorbances


           When two absorbing species are present in solution, the absorbance value measured will represent both species

      A  =  A1  +  A2

      However, measuring absorbance at two different wavelengths can yeild the exact concentration of each component


      A1     =   e1   b   C1   +    e2'  b   C2                      at l1


      A2     =    e1'  b   C1   +    e2  b   C2                      at l2


       e1   ,    e2     ,    e1'   ,  and   e2'  are constants and can be determined experimentally.


      Only C1  and C2  are not known which can be determined by solving the two equations.


Determination of a Ligand to Metal ratio


           UV-Vis spectroscopy is very useful in determining the ratio between a ligand and a metal.


      Mn+  +  m L =  (MLm)n+


      Usually, complexation of ligands with metals result in different spectroscopic characteristics for both.  The most pronounced situation is the formation of a colored complex.  Two widely used experimental methods will be discussed:


      a. The Method of Continuous Variation (Job's Method)


           In this method, the mole fraction of either the metal or the ligand is plotted against absorbance.  This yields a result similar to that shown in Figure 4.



      Figure 4 : A plot of absorbance versus mole fraction.



                                       mol Mn+

      Mole Fraction =   ______________

                               mol Mn+  +  mol L


      Tangents are drawn on both side of the maximum obtained and a perpendicular line is drawn to the axis representing the mole fraction.  This gives the mole fraction of the metal.  If the value is  0.5  then it is a 1:1 complex and if it is  0.33  then it is a 1:2  complex, etc.

      The method of continuous variation is excellent for complexes that are 1:1 but if the ratio is more than 1:2 there will be some considerable uncertainty and the mole ratio method is preferred.


      b. Mole Ratio method


           The concentration of the metal ion is usually kept constant and a variable amount of the complexing agent is added.  The mole ratio of the metal ion to the ligand is plotted versus absorbance and a result as shown in Figure 5 is obtained



                                       mol L/mol Mn+


      Figure 5: A plot of absorbance versus mole ratio.



      Tangents are drawn and a perpendicular line is drawn to the mole ratio axis showing the exact ratio.





           In this section, only a brief description of instrumental features will be mentioned.  This is important since you may be required to perform some experiments in UV-Vis spectroscopy without enough background.

      Two types of instruments are available according to the wavelength selector used.


      a. Filter Photometer


           This uses filters for the selection of working wavelengths.  Photometers are cheap machines that are widely used in most primitive analytical laboratories.  The optical system and instrumental components can be represented by Figure 6.




      Figure 6:  Schematic diagram of a photometer


           As can be seen from the figure, light is emitted from the source passing through a suitable filter for wavelength selection.  Part of the light at the selected wavelength is absorbed by the sample and the transmitted light hits the phototube detector resulting in a signal that is displayed by the instrument as absorbance.


      b. Dispersive Spectrophotometers


           These use either prisms or gratings for wavelength selection.  Prisms and gratings are excellent wavelength selectors where a very narrow band of light at specific wavelength can be chosen especially with good gratings.  Dispersive instruments are divided into two types:


      1.  Single Beam Spectrophotometers


           This is similar to the photometer design but the wavelength selector is either a prism or grating instead of the filter.  Usually, single beam instruments are of moderate price and require adjustment to zero using a blank before sample measurement.  As the instrument is kept in the operational mode, multiple zero adjustments should be undertaken because there is always some drift in response with time.


      2.  Double Beam Spectrophotometers


           These incorporate places for two cells one for the blank and the other for the sample.  The instrument automatically subtracts the absorbance of the blank or reference from that of the sample.  The different optical components of the instrument can be seen in Figure  7 .



      Figure 7:  Schematic diagram of a double beam spectrophotometer: L,source; G, grating; M, mirror; C, chopper; S, sample; R, reference;    and D, detector.


           The chopper (C) splits the incident beam into two halves, one passes through the sample and the other passes through the reference.  The detector automatically records the difference which is displayed as absorbance.




           The most commonly used sources are deuterium lamps in the ultraviolet region and tungsten - halogen lamps in the visible region.  Make sure not to look at the deuterium lamp while in the operational mode since UV light is damaging to your eyes.




           Remember that glass absorbs UV light, therefore make sure to use quartz cells when working in the UV region.  Glass cells are adequate for measurement of absorbance in the visible region while quartz cells are adequate through the whole UV-Vis range.


      Routine Methodology in Spectrophotometric Analysis


           The first step of an analytical procedure in UV-Vis spectroscopy is to find the wavelength that yeilds maximum absorbance.  This is done by scanning through the UV or Vis range, depending on the characteristics of the absorbing species.  The spectrum is plotted with absorbance on the Y-axis and the wavelength on the X-axis.  Then the wavelength that yeilds maximum absorbance is chosen for further work.  This also gives maximum molar absorptivity.

           When the problem involves the determination of an unknown analyte concentration, standard analyte is used to construct a calibration curve at the preselected wavelength and the unknown absorbance is measured which can be correlated with concentration from the curve.








Experiment 1:  Plotting the Absorption Spectrum of Iron (II) Complex with 1,10- Phenanthroline




           The first step of an analytical spectrophotometric procedure for quantitative determination of analytes is to find the wavelength at which the analyte complex has maximum absorption. At this wavelength, the molar absorptivity is a maximum and precision is greater. This allows for more precise and sensitive determinations.  Iron (II) forms an orange-red complex with 1,10-phenanthroline (phen) in the pH range from 2-9.  The absorption spectrum of this colored complex is to be determined.




n Fe2+  +  m phen = Fen(phen)m2+  (orange-red)





      1. A spectrophotometer or a photometer.

      2. Sample cells.

      3. pH Meter.


Chemicals and Reagents


      a. Provided


      1. 0.1 M Acetate buffer, pH 4.0.

      2. 0.5 M Hydroxylamine hydrochloride.

      3. Predried and desiccated Fe(NH4)2 (SO4)2 6H2O.

      4. Predried and desiccated 1,10- phenanthroline.


      b. Need Preparation


      1. 1 L of 5x10-4 M Fe2+ solution.

      2. 1 L of 5x10-4 M phen solution.

      Note:  Solutions 1 and 2 will be used for upcoming experiments, therefore make sure to save them.




      1. Accurately transfer 5 mL portion of Fe2+ standard solution and 5 mL of the acetate solution into a 25 mL volumetric flask.

      2. Pipet exactly 2 mL of hydroxylamine hydrochloride solution provided followed by a 5 mL portion of the standard phen solution.

      3. Adjust the volume to the mark using dustilled water and allow to stand for 10 min.

      4. Repeat the same procedure from 1 to 3 but with no addition of Fe2+ in step 1. This is  your blank solution.

      5. Follow instructions for operation of the photometer or spectrophotometer provided and adjust it to zero according to manufacturer's instructions.  Discussion below will be related to a single beam photometer.

      6. Place the blank in one of the absorption cells provided and adjust the absorbance reading to zero  at 450 nm.

      7. Remove the blank cell and place the sample cell in its proper place and measure absorbance at 450 nm. Record your results.

      8.Repeat steps 6 and 7 each time you perform a measurement at a new wavelength. Record your results for wavelengths from 450 to 550 nm in 5 nm intervals.

      7. Plot your results as an x-y graph with absorbance on the y axis and wavelength on the x axis and draw the absorption spectrum.

      8. Find the wavelength that results in maximum absorption. This is the lmax which you will use for any future quantitative work on this complex.


Experiment 2:  Determination of Iron in aqueous Samples




           Iron present in aqueous samples can be determined  spectrophotometrically by complexation with a suitable complexing agent. The absorbance of the metal-ligand complex is usually measured in the visible region and is related to metal ion concentration.

           Colorimetric determination of iron can be done using several known complexing agents. Among the routinely used is 1,10-phenanthroline (phen) which reacts with Fe2+ to form an orange-red complex. Therefore, the first step involves the reduction of any Fe3+ present to Fe2+ using hydroxylamine hydrochloride.

      The procedure depends on the construction of a calibration curve from standard Fe2+, followed by measurement of the unknown Fe2+ concentration from the curve.




      n Fe2+  +  m phen  =  Fen (phen)m2+





      1. A photometer or spectrophotometer.

      2. Sample  cells.

      3. pH Meter.




Chemicals and Reagents


      a. Provided


      1. Stock phen solution (from previous experiment).

      2. 0.5 M Hydroxylamine hydrochloride.

      3. 0.1 M Acetate buffer, pH 4.0.

      4. Standard Fe2+ solution (from previous experiment.


      b. Need Preparation

      None, except 1 and 4 if you do not have them ready. Procedure for preparation of solutions 1 and 4 is described in previous experiment.




      1. Accurately transfer 0,1,2,3,4,5 and 6 mL of standard Fe2+ solution (solution 4) into separate 50 mL measuring flask.

      2. To each solution in step 1 add 5 mL of acetate buffer followed by 2mL of hydroxylamine hydrochloride.

      3. Transfer exactly 20 mL portions of the stock phen solution to each measuring flask and complete to mark with distilled water.

      4. Repeat steps 1-3 replacing the unknown (5 mL) for standard Fe2+, three times.

      5. Allow 10 minutes for color development and measure absorbance of each solution at the wavelength obtained from previous experiment using the solution with zero mL Fe2+ as your blank.

      7. Find the concentration of the unknown Fe2+ from the calibration curve and report your results as ppm Fe.

      8. Find the relative standard deviation of the unknown Fe2+ concentration in ppm.


      9. Find the molar absorptivity of the Fen(phen)m2+ complex from the slope of the curve.


Experiment 3 :  Determination of the Mole Ratio of Iron : 1,10- phenanthroline in the complex




           The mole ratio of a metal to a ligand in a metal-ligand complex can be determined. In the complex, Fen(phen)m2+, the coefficients n and m can be calculated by several methods. Two very important, simple methods, will be applied in this experiment. The method of continuous variation and the mole ratio methods are widely used with advantage of the first if the ratio is 1:1 .  The second method is excellent and better than the first when the mole ratio is greater than 1:2.




      n Fe2+  +  m phen  =  Fen (phen)m2+



Chemicals and Reagents


      a. Provided


      1. 0.1 M Acetate buffer, pH4.0.

      2. 0.5 M Hydroxylamine hydrochloride.

      3. Predried and desiccated 1,10-phenanthroline.

      4. Predried and desiccated Fe(NH4)2 (SO4)2 .6H2O


      b. Need Preparation


      1. 5x10-4 M 1,10-phenanthroline solution.

      2. 5x10-4 M Fe2+ solution.

      Note: Use solutions saved from previous experiments.




I. Application of the Method of Continuous Variation


      1. Accurately transfer 0,1,2,3,4,5,6,7,8,9,and 10 mL of the Fe2+ solution into  mL measuring flasks and mark each flask.

      2. To each flask add 5 mL of the acetate buffer provided followed by 2 mL of hydroxylamine hydrochloride.

      3. Pipette 10,9,8,7,6,5,4,3,2,1,and 0 mL of the phen solution into the flasks in step 1, respectively. Each flask should now contain 17 mL of the mixture.

      4. Dilute to the mark using distilled water and allow the solutions to stand for 10 min.

      5. Record the absorbance of each solution using water as a reference.

      6. Plot the absorbance of the solutions versus mole fraction of Fe2+ and find the stoichiometry of the reaction.


II. Application of the Mole-Ratio Method


      This involves addition of fixed amounts of the metal ion to each solution and changing the number of moles of the complexing agent.

      1. Accurately transfer 4 mL of the standard Fe2+ solution into nine 50 mL volumetric flasks and mark your flasks.

      2. To each flask add 5 mL of the acetate buffer followed by 2 mL of the hydroxylamine solution.

      3. Pipet exactly 1,3,6,8,10,12,14, 17 and 20 mL of the phen solution and complete to the mark with distilled water.

      4. Allow 10 min before measurement of absorbance at   lmax.

      5. Plot the absorbance of the different solutions versus the mole ratio of phen to Fe2+ and identify the stoichiometry of the complex from the resulting curve.




      It is hoped that, by now, you will have an appreciation of some basic uses of UV-Vis spectroscopy. You should be able to find lmax, concentration of an unknown from a calibration plot, the molar absorptivity as well as the stoichiometry of any colored metal-ligand reaction.


      To have a comprehensive appreciation you should study the different variables that affect the reaction system , like pH, reactants concentration, temperature, ...etc.