environmental geochemistry course lab report

Environmental Geochemistry Course Lab Report

Prashant MahendranDate of Submission: 1/6/15

Overall Introduction

The environmental geochemistry lab course is a three-week lab course involving experiments that are readily applicable to modern day geoscience. Specific lab techniques for each experiment are introduced throughout the course. The 5 experiments performed are as follows:

1. Arsenic Detection and speciation in drinking water and electrochemical detection2. pH, buffers and alkalinity3. Gibbs energy changes during Cobalt complexation4. Determination of Ksp, G, H and S for the dissolution of calciumΔ Δ Δ hydroxide in water5. Oxidation Kinetics of Fe2+ in aqueous solutions

These experiments are performed in groups of 3 with the supervision of a teaching assistant, before lab reports are written and handed in the next day.

One of the experiments is performed in more detail as an independent experiment; I chose Oxidation Kinetics of Fe2+ in aqueous solutions and used more pH values and solutions of varying ionic content.

Problem sets

1.3 Chemical measurements

4. a) Molarity of NaCl in the ocean:2.7g per 100ml- 27g per litre.No. of moles = mass/molar mass= 27/(23.0+35.5) = 27/58.5 = 0.46 M

b) Mass of MgCl2

Mass= no.of moles * molar mass = 0.054*(24.3 + 2(35.5)) = 0.054*95.3 = 5.1462g in 1 litre. In 25ml- 0.025*5.1462= 0.129g

5. Molarity= moles of HCl/volume of solution

Mass of solution= density*volume= 1.19*1000=1190g in 1 litreMass of HCl per litre=1190*0.37=440g/lMolarity=Moles HCl per litre=mass/molar mass=440/36.5=12.1 M

Molality= moles of HCl/mass of waterIn 100g solution- 37g HCl, 63g H2O. Moles HCl= 37/36.5= 1.01 molMolality= 1.01/0.063 = 16.1 m

6. 8 mM = 0.008 MSo 0.008 mol of Copper(II) sulphate pentahydrate per litre0.008*0.5= 0.004 mol of Copper(II) sulphate pentahydrateMass= no.of moles * molar mass = 0.004* 249.68= 0.999g

7. a) 0.1 M HCl in 1lMoles of HCl required= 0.1 Volume= 0.1/12.1=0.00826 litres = 8.26ml conc. HCl.

b) 0.25M NH3 so 0.25*0.5=0.125 mol NH3 in 500mlConcentrated ammonium hydroxide- 1 litre is 0.899*1000=899 g28% wt so 899*0.28= 251.72g in 1 litreMoles= mass/molar mass= 251.72/(14+3)= 14.81mol in 1 litre= 14.8 MVolume= 0.125/14.8=0.00844 l= 8.44ml

1.4 Tools of the trade and Experimental error8.

9. ±0.04

1.5 Statistics10.

11.

Equation:y±0.006 = (m±0.0002)x + (b±0.003)

1.6 Quality assurance and calibration methods12. Standard addition equation:Concentration of analyte in initial solution/concentration of analyte plus standard in final solution= signal from initial solution/signal from final solution[Na]i/[S]f+[Na]f=Ix/Is+x

Where [S]f=[S]i(Vs/V) where the quotient is the dilution factor.[Na]f=[Na]i(V0/V) Vs is added volume, V0 is initial volume and V is final volume.[Na]i/([Na]i(0.95)+2.08(0.05))=4.27/7.98[Na]i=0.113M

13.

Arsenic Detection and speciation in drinking water and electrochemical detection

Group 5: Robert Yates, Prashant Mahendran, Matthew Kirby

Submitted 12/5/15Edited version: Prashant Mahendran

AbstractThis report aims to look at detection of trace amounts of arsenic in a variety of water sources. The technique used is differential pulse anodic stripping voltammetry (DPASV). It was found that Volvic and Highland Spring drinking waters did not exceed World Health Organisation arsenic limits. IntroductionArsenic contamination in drinking water is a serious problem in more than 70 countries affecting over 137 million people (Smedley and Kinniburgh, 2002). The World Health Organisation has set a maximum safe limit of 10ppb (Bang et al., 2005). However, due to extremely low concentrations, difficulties arise from detecting such low concentrations of Arsenic contamination. This requires very sensitive equipment in order to detect such low quantities. One method of detecting very low concentrations is Anodic Stripping Voltammetry.Anodic Stripping Voltammetry is an electroanalytical method where information about the analyte is obtained by measuring the current (signal) as the potential is varied. There are several ways to measure the potential, in this case differential pulse method was used.

Anodic Stripping Voltammetry involves the use of three electrodes:Working electrode – This will make contact with the analyte (arsenic particles) and applies the desired potential in a controlled way, facilitating the transfer of charge to and from the analyte. In this case a gold microwire electrode is used. At the surface of the gold electrode, the arsenic interacts with the gold and is

reduced to As(0) and sticks to the surface. In stripping voltammetry, after a certain amount of time the arsenic will

be oxidised to As(III) or As(V), and therefore will be released off the surface. These redox reactions generate the signal.Reference electrode – Measure the potential at the working electrode from a stable reference point. In this case a silver electrode is used which contains sodium nitrate and potassium chloride solution.

Fig 1 Diagram illustrating Anodic Stripping Voltammeetry (after Jahandari

et al, 2013)

Auxiliary electrode – Balances the charge from the working electrode. In this case an Iridium wire is used.The aim of this report is to calculate the Arsenic concentrations, through the use of anodic voltammetry, in four different samples of unknown concentration.The method of standard addition will be used in place of a calibration curve. This involves spiking the solution with known concentrations of arsenic and extrapolating the graph of signal against added arsenic to find the original arsenic concentration. This technique is repeated for each sample, and avoids the matrix effect (other ions in the sample affecting the signal, making a calibration curve unreliable for different samples). Experimental sectionMethod

1. Condition the gold electrode of the voltammeter by running a cyclic voltammetric scan with 0.5M H2S04. The program was set to use a conditioning potential of -2.5V, a duration time of 30 seconds and an equilibrium time of 2 seconds.

2. We then added 500 µL of HCl (6M) to a 30 mL sample of Milli-Q water to produce an overall concentration of 0.1M of HCl. This is then used as a reagent blank.

3. We then mixed the solution for about 30 seconds by inserting a small magnet and using a magnetic stirrer.

4. Once the solution had been mixed, the program was run 5 times. The following parameters were used

a. Purge time = 0b. Conditioning potential = 0.7V; Duration = 2sc. Depositional potential = -1V; Duration = 10s (analytical) 1s

(background check)d. Equilibration time = 1s

5. Spike the solution with 80µL of Arsenic and repeat steps 3 and 46. Repeat step 5 twice.7. Run the background subtraction (BGS) programme to remove the matrix

effects from the detected peaks.8. Using the peak derivatives from the arsenic peaks (the last 3 out of 5

scans) from each of the graphs the peak derivative can be plotted against added concentration. From these points a line of best fit was made and extrapolated to find the arsenic concentration (using the standard addition method).

9. Compare this result to find the known value of the added solution (since pure Milli-Q water was used). This gives us the recovery of the method (accuracy).

10. Rinse the electrodes and cell and repeat steps 2-8 using each of the 4 samples of unknown arsenic concentration (instead of Milli-Q water). The standard addition method will correct for the matrix effect problem. (We only had time to do 2 samples as the electrode malfunctioned).

ResultsA solution of known As concentration was spiked with known amounts of As to determine the recovery of the system. Using the equation y=mx+c using original concentration was found to be 1.89ppb.This is not the ~2.7ppb used; the recovery of the system was 72%. However the linearity is quite high/close to 1

0 1 2 3 4 5 6

0.00E+00

1.00E-07

2.00E-07

3.00E-07

4.00E-07

5.00E-07

6.00E-07

7.00E-07

8.00E-07

9.00E-07

1.00E-06

R² = 0.999427745406392f(x) = 1.23740695833333E-07 x + 2.33294444444448E-07

As added (ppb)

Added AsV ppb

Deriv

ative

Fig 2 Standard Addition graph to determine As concentration from spiking milli-q water with known As concentrations.

Vo(ml)= Volume added

(Vs,mL)=

Peak derivative

x Axis function([As]*Vs/Vo)

y-axis function= peakderivative*(V/Vo

)30 0 1.94E-07 0.00 1.94E-07

[As] std(ppb)=

0.08 4.68E-07 2.67 4.70E-07

1000 0.16 7.46E-07 5.33 7.50E-07Table 1 Results for Highland Spring water

0.00 1.00 2.00 3.00 4.00 5.00 6.000.00E+00

1.00E-07

2.00E-07

3.00E-07

4.00E-07

5.00E-07

6.00E-07

7.00E-07

8.00E-07

f(x) = 1.04390116666667E-07 x + 1.92802859259258E-07R² = 0.999975400609171

Spring Water

Added Concentration(ppb)

Pea

k D

eriv

ativ

e

Fig 3 Determination of As concentration in Highland Spring waterThe graph shows a concentration of ~ 2 ppb. While using the equation y = mx +c an exact value of 1.85 ppb was calculated.

Vo(ml)= Volume added (Vs,mL)=

Peak derivative x Axis function([As]*V

s/Vo)

y-axis function= peakderivative*(V/Vo

)30 0 2.07E-07 0.00 2.07E-07

[As] std(ppb)= 0.08 6.05E-07 2.67 6.06E-071000 0.16 9.46E-07 5.33 9.51E-07Table 2 Results for Volvic water

0.00 1.00 2.00 3.00 4.00 5.00 6.000.00E+00

1.00E-07

2.00E-07

3.00E-07

4.00E-07

5.00E-07

6.00E-07

7.00E-07

8.00E-07

9.00E-07

1.00E-06

f(x) = 1.39483133333333E-07 x + 2.1597475555556E-07R² = 0.998164807353093

Volvic Water

Added Concentration(ppb)

Pea

k D

eriv

ativ

e

Fig 4 Determination of As concentration in Volvic waterThe graph shows a concentration of ~ 1.5 ppb. While using the equation y = mx + c an exact value of 1.55 ppb was calculated.Discussion

Quality control: the standard addition procedure was performed by spiking a solution of known arsenic concentration, resulting in a recovery of 72% which gives us concentrations of only a moderate accuracy. However it can be concluded that these results show levels of Arsenic concentrations in drinking water below 2ppb, well below the WHO maximum limit for safe consumption of 10ppb. It can also be seen that the level of Arsenic in the spring water is approximately ~0.3ppb higher than the volcanic water. This initially appears to be a surprising result, considering that the proximity of the source to volcanic activity would likely increase Arsenic concentrations. However, this result could be due to a difference in the method of filtration or a general difference in the geology of the area (clay soils may adsorb arsenic and remove it from solution due to a negative surface charge of the clay) or residence time of the water underground. Unfortunately, due to the fragility of the extremely thin gold microwire electrode the experiment was not as successful as initially planned for the timescale. This was necessary for the level of accuracy required to measure trace elements. The arsenic was not recorded for the Thames water and Serpentine water samples to allow more comparison.

ConclusionWe have found that the drinking water falls below WHO limits, however there is some variation of concentration between the two bottle water types which could be due to a variety of factors such as the local geology, and the companies filtering techniques. In the future the experiment could be run for waters found in natural sources to test the safety of these for drinking. These results can then be compared with the drinking water. It would also be good to test how much As(III) there is compared to total As, as As(III) is 600x more dangerous than As(V).

ReferencesBang, S., M. Patel, L. Lippincott, and X. Meng (2005), Removal of arsenic from groundwater by granular titanium dioxide adsorbent, Chemosphere, 60, 389-397.Smedley, P. L., and D. G. Kinniburgh (2002), A review of the source, behaviour and distribution of arsenic in natural waters, Appl. Geochem., 17(5), 517-568.Jahandari, S., Taher, M.A., Fazelirad, H., Sheikhshoai, I. (2013), Anodic stripping voltammetry of silver(I) using a carbon paste electrode modified with multi-walled carbon nanotubes, Microchimica Acta, 180(5-6), 347-354

pH, buffers and alkalinity

Group 5: Robert Yates, Prashant Mahendran, Matthew KirbySubmitted 11/05/2015

Edited version: Prashant MahendranAbstractTitration experiments were carried out to determine the buffer action for both an acid (phosphoric acid) and a base (the carbonate system in the Serpentine). Titration against phosphoric acid showed a stepwise buffering action due to the polyprotic nature of phosphoric acid. Maximum buffer capacities were at pH 2.66 and pH 7.03 and pKa1=2.75, pKa2=7.1. Results obtained showed a significant buffering capacity in the lake studied and thus a low sensitivity to acid influx. The total alkalinity found was 146mg/L.IntroductionA buffer solution resists changes in pH(potential hydrogen, -log[H+]) when small amounts of acid or base are added to it. For a general dissociation reaction of an acid, the equation is:HA= A- +H+ (1)Henderson-Hasselback equation: pH= pK- log [HA]/[A-] (2)where pK is –log(Keq). Keq is the equilibrium constant. pK, or pKa is specific for a weak acid. The best buffering action is when [HA]=[A-], when pH=pK. pK can be determined from the titration experiment by adding a strong base to a weak acid or strong acid to a weak base and finding the midpoint of plateaus on the titration curves (pH against added acid/base).Alkalinity is the ability of a solution to neutralise the addition of acids. In the carbonate system, the total alkalinity is defined asAlk= [HCO3

-] + [CO32-] (3)

The aim of the experiment is to find the pKas of phosphoric acid by titrating it with sodium hydroxide, and then titrate hydrochloric acid into water from the Serpentine (a lake) to measure the buffering capacity (alkalinity).

Experimental SectionMethod:Experiment 1

1. 0.257g solid phosphoric acid was weighed and dissolved into 400ml distilled water to form a 0.00525M solution (0.005M was the aim, however it was too difficult to measure the exact mass required for this). This was then poured in a 500ml beaker.

2. 20ml of 1M NaOH solution was mixed with 80ml milli-Q water to make 100ml of 0.2M NaOH solution.

3. A magnet was placed into the beaker and a magnetic stirrer was used to mix the solution. This was left on constantly throughout the experiment. The pH electrode was clamped in place in the beaker and allowed to acclimatise.

4. 1 ml aliquots of 0.2 M sodium hydroxide were titrated into the beaker from a burette, taking measurements from the pH electrode 30 seconds after adding each of the aliquots.

5. This was continued until pH 11.

6. This data was plotted on a scatter chart with volume of NaOH added on the x axis and pH on the y axis. The plateaus of this titration curve will be used to estimate two of the pK values of the phosphoric acid.

7. Using the gradient of the graph and converting the units into mole equivalents, buffer capacity ( ) was calculated and plotted on a graph βagainst pH. The maximum values occur when pH=pK. β

Experiment 21. Water from the serpentine lake was filtered to remove undissolved

particles.2. 50ml of the water was added to a 100ml beaker.3. A pH meter was placed in the beaker in the same way as experiment one.4. 0.1 ml aliquots of 0.1 M HCl were periodically added to the beaker with a

pipette.5. The pH reading was recorded after each aliquot was added.6. The pH was plotted against the change in volume.7. The graph was used to find the equivalence points, and the buffer areas

were observed (plateaus). 8. Using derivatives and the Gran plot, end points for the reaction were

determined.

ResultsNaOH with H3PO4

0 5 10 15 20 250

1

2

3

4

5

6

7

8

9

10

11

NaOH with H3PO4

Volume of NaOH added (ml)

pH

Fig 1 graph showing titration of NaOH with H3PO4 pKa1=2.75, pKa2=7.1

Volume of alkali pH Gradient0 2.49 dpH/dV dV/dpH dm/dpH moles of acid β1 2.55 0.06 16.67 3.33 0.0021 1587

2 2.62 0.07 14.29 2.86 0.0021 13613 2.66 0.04 25.00 5.00 0.0021 23814 2.76 0.1 10.00 2.00 0.0021 9525 2.85 0.09 11.11 2.22 0.0021 10586 2.96 0.11 9.09 1.82 0.0021 8667 3.16 0.2 5.00 1.00 0.0021 4768 3.42 0.26 3.85 0.77 0.0021 3669 4.14 0.72 1.39 0.28 0.0021 132

10 6.04 1.9 0.53 0.11 0.0021 5011 6.43 0.39 2.56 0.51 0.0021 24412 6.66 0.23 4.35 0.87 0.0021 41413 6.88 0.22 4.55 0.91 0.0021 43314 7.03 0.15 6.67 1.33 0.0021 63515 7.23 0.2 5.00 1.00 0.0021 47616 7.45 0.22 4.55 0.91 0.0021 43317 7.70 0.25 4.00 0.80 0.0021 38118 8.44 0.74 1.35 0.27 0.0021 12919 10.17 1.73 0.58 0.12 0.0021 5520 10.65 0.48 2.08 0.42 0.0021 19821 10.90 0.25 4.00 0.80 0.0021 38122 11.00 0.1 10.00 2.00 0.0021 952

Table 1 Data for calculation of buffer capacity of the NaOH. Highlighted values correspond to equivalence points.

2 3 4 5 6 7 8 9 10 11 120

500

1000

1500

2000

2500

Buffer Capacity

pH

β

Fig 2 Graph showing buffer capacity of NaOH with H3PO4.Max. buffer capacities at pH 2.66 and pH 7.03 (highlighted on table).

HCl titrated against Serpentine water

Fig 3 Graph of pH against volume of HCl added to Serpentine Lake water, with corresponding graphs of first and second derivatives. The graph of pH against volume of HCl plateaus at pH 3 and pH 6.5, with the larger plateau at 6.5. Using the second derivative graph, the end points are at 1.3ml and 1.9ml. HCl added(ml) pH 10^-pH vb*10^-pH

0 8.33 4.68E-09 00.1 7.8 1.58E-08 1.58E-090.2 7.55 2.82E-08 5.64E-090.3 7.25 5.62E-08 1.69E-080.4 7.09 8.13E-08 3.25E-080.5 6.95 1.12E-07 5.61E-080.6 6.78 1.66E-07 9.96E-080.7 6.67 2.14E-07 1.50E-070.8 6.56 2.75E-07 2.20E-070.9 6.48 3.31E-07 2.98E-07

1 6.39 4.07E-07 4.07E-071.1 6.29 5.13E-07 5.64E-071.2 6.17 6.76E-07 8.11E-071.3 6.11 7.76E-07 1.01E-061.4 6.02 9.55E-07 1.34E-061.5 5.93 1.17E-06 1.76E-061.6 5.82 1.51E-06 2.42E-061.7 5.59 2.57E-06 4.37E-061.8 5.35 4.47E-06 8.04E-061.9 4.92 1.20E-05 2.28E-05

2 4.08 8.32E-05 1.66E-042.1 3.65 2.24E-04 4.70E-042.2 3.42 3.80E-04 8.36E-042.3 3.37 4.27E-04 9.81E-042.4 3.16 6.92E-04 1.66E-032.5 3.07 8.51E-04 2.13E-032.6 2.97 1.07E-03 2.79E-032.7 2.89 1.29E-03 3.48E-032.8 2.82 1.51E-03 4.24E-03

Table 2 Data from titration of HCl with Serpentine water used to create a Gran plot.

0 0.5 1 1.5 2 2.5 30.00E+00

5.00E-04

1.00E-03

1.50E-03

2.00E-03

2.50E-03

3.00E-03

3.50E-03

4.00E-03

4.50E-03

Gran Plot

Volume

Vo

lum

e*

10

^-p

H

Fig 4 Gran plot for titration of HCl with Serpentine water.

Gran PlotThe x axis intercept on the gran plot shows an end point at 2.1ml. This end point corresponds to the point at 1.9ml in the second derivative plot. Alkalinity=

(Concentration of added acid*volume of added acid)/Volume of initial sample=(0.1*0.002)/0.050=0.004M=146ppm=146mg/LDiscussionExperiment 1:The stepwise buffering action showed the polyprotic nature of phosphoric acid. Only 2 dissociation steps were observed as the third step is too alkaline for the pH meter to record accurately. H3PO4 = H+ + H2PO4

- (pKa1)H2PO4

- = H+ + HPO42- (pKa2)

HPO42- = H+ + PO4

3- (pKa3, not recorded)The titration curve produced clear plateaus to calculate the pKa values from (pKa1=2.75, pKa2=7.1). The graph plotted of buffer capacity against pH showed 2 clear peaks which corresponded to these pKa values (max. buffer capacities at pH 2.66 and pH 7.03).Experiment 2:The main buffering system in the Serpentine water is the carbonate system.Equation 1: H2CO3 = HCO3- + H+Equation 2: HCO3- = CO3- + H+The equations move to the left upon addition of acid/H+ (Le Chatelier’s principle). The literature pKa value for equation 1 is pka1= 3.6 for H2CO3 only and pka1=6.3 including dissolved CO2 (Silberberg, 2009). For equation 2 pKa2=10.32 (Silberberg, 2009).Since the serpentine water is an open system in contact with atmospheric CO2, The dissolution of CO2 replenishes H2CO3, resulting in more of the dissociation occurring over time, hence the very large plateau indicating intense buffer action of the system.The graph of pH against volume of HCl plateaus at pH 3 and pH 6.5, with the larger plateau at 6.5. These correspond to pKas of the buffer solution of 3 and 6.5, with more buffering capacity at 6.5. The literature values for the pKa1 of carbonic acid (H2CO3) are 3.6 for H2CO3 only and 6.3 including dissolved CO2 (Silberberg, 2009). The total alkalinity found was 146mg/L. Therefore, the buffering action can be assumed to primarily originate from the carbonate system. Typical alkalinities for freshwater lakes are 20-200mg/L CaCO3 (Hudson, 1998), so the Serpentine has a high buffering capacity and thus is not very sensitive to acidification (for example from acid rain). ConclusionsIn conclusion, titration experiments on a variety of scales enable us to calculate specific pKas and thus gain an understanding of buffering action and total carbonate alkalinity in natural water systems. This can therefore allow environmental agencies to monitor specific freshwater systems for their vulnerabilities to acid influx, particularly due to industrial pollution or acid mine drainage from nearby sources. ReferencesHudson, H. (1998), Lake notes, 4pp., Illinois environmental protection agencySilberberg, M. S. (2009), Chemistry, 5th edition, McGraw-Hill

Gibbs energy changes during Cobalt complexationGroup 5: Robert Yates, Prashant Mahendran, Matthew Kirby

Date of submission: 12/05/2015

Edited Version: Prashant MahendranAbstractGibbs free energy values were calculated for cobalt complexation between the [Co(H20)6] and [CoCl4] species using data from UV/Vis molecular absorption spectroscopy. Enthalpy and entropy were then derived by repeating the experiment at 20 ˚C, 30 ˚C, 40 ˚C and 50˚C and plotting a graph of G˚ against Δtemperature. The reaction is endothermic with an increase in disorder. Assuming that enthalpy and entropy remain constant over small temperature ranges these were derived to be 22.4kJ/mol and 79.9J/mol/K respectively.IntroductionThe aim of this experiment is to determine the Gibbs free energy change for Cobalt complexation between the aqueous [Co(H20)6] form and the [CoCl4] chloride form. To achieve this aim, we added a large quantity of Cl- to the [Co(H20)6] solution containing a red octahedral complex and a blue tetrahedral complex is produced. When the solution temperature is modified (20 ˚C, 30 ˚C, 40 ˚C and 50˚C), the equilibrium constant, Keq, of the complexation reaction is shifted thus altering the proportion of the two coloured forms correspondingly. Then a graph of G˚ against temperature (in Kelvin) can be plotted to find the Δ

S˚ and H˚.Δ ΔThe absorption of a specific complex in the UV/Visible light range is determined by the energy difference between d-orbitals in the transition metal ion. The degenerate (energetically equal) orbitals are split by ligands. The splitting is determined by the arrangement of the ligands geometrically around the transition metal ion, the ligands present and the oxidation state of the metal. The light absorption excites the electrons and promotes them into the higher energy d-orbitals. The energy of the light absorbed determines the colour of the complex. Higher energy of light absorbed corresponds to lower wavelength, i.e. a higher energy difference between d orbitals in [Co(H2O)6] compared to CoCl4 results in lower wavelength(bluer) light being absorbed by [Co(H2O)6] resulting in its red colour.The Beer-Lambert law used to calculate concentration:Absorption= .c.L (1)εWhere =molar extinction coefficient (L/mol/cm), c= concentration of species ε(M) and L=path length (cuvette size, cm)The equation for the reaction is:Co(H2O)6+4Cl-= CoCl4 +6H2O (2)Equation for equilibrium constant Keq:Keq=[CoCl4]/[Co(H2O)6]*[Cl-]4 (3)Calculating gibbs free energy from Keq:ΔG= RTln(Keq) (4)Where R=gas constant and G=gibbs free energyEquation for Gibbs free energy: ΔG= ΔH-TΔS (5)Where G=gibbs free energy(in kJ/mol), H=enthalpy (in kJ/mol), T=Temperature in Kelvin and S=entropy(in KJ/mol/K)The key techniques being learnt were:

Producing samples for UV/Vis spectroscopy in cuvettes. Calibration of the equipment using instrumental blanks (milli Q water). The importance of experimental blanks.

Interpreting graphs of results plotted on the computer by the machine (absorption spectra over a range of wavelengths).

Calculating concentrations of species from absorptions using the Beer-Lambert Law.

Using equilibrium equations to derive equilibrium constants.

Experimental sectionMethod

1. Using the molar mass labelled on the container of CoCl2.6H20 the mass required to form a 50ml 0.09M solution in water was found to be 1.07g.

2. This was weighed out on a weighing boat and mixed with pure water in a 50ml volumetric flask to dissolve it.

3. This solution was pipetted into 4 sets of three plastic containers in 2.5ml aliquots.

4. 4 blank solutions were also made with 2.5ml aliquots of pure water.

5. All of the solutions were then diluted with 2.5ml of concentrated (12.1M) HCL.

6. The solutions were then equilibrated to 24 ˚C in a water bath. 7. A baseline was created on the UV/Vis spectrometer by placing a

cuvette of pure water into the machine as an instrumental blank. This was for calibration.

8. The first set of 4 plastic containers (3 containing cobalt + HCL and one blank of just H20 + HCL) were pipetted into cuvettes (Approximately 1ml).

9. The experimental blank of just H20 + HCL was first run through the machine to produce a spectrum with no peaks.

10. The three solutions at a temperature of approximately 24 ˚C were then run through the machine.

11. Absorption readings (and blank readings) were recorded at 2 peaks for each of the 3 solutions: 522nm(absorption for Co(H2O)6) and 690nm(CoCl4 absorption).

12. This was repeated for a set of 3 containers and 1 blank at 30˚C, 41˚C and 51˚C.

13. The Beer-Lambert law was used to calculated the concentrations of the species present in the solution at each temperature (from an average of the 3 containers).

14. These concentrations were then used to determine the equilibrium constants Keq at each temperature.

15. The equation G˚=-RTln(Keq) was used to find G˚ values for the Δ Δreaction, and these were plotted on a graph G˚ against Δtemperature (in Kelvin) can be plotted to find the S˚ and H˚.Δ Δ

Results

The Peak at 522nm did not vary significantly between experiments so we have only considered the peak at 690nm and determined [Co(H2O)6] concentration using the total Cobalt concentration(calculated initially as 0.09M) minus the CoCl4 concentration determined from the absorbance results.

Absorption(A) conc. Co(H2O)6

Temp(C) Blank 1 2 3 Average Av - blank

abs/ε.l

24 -0.003 0.392 0.406 0.404 0.401 0.404 0.08830 0.025 0.433 0.411 0.410 0.418 0.393 0.08541 0.033 0.419 0.422 0.413 0.418 0.385 0.08351 -0.003 0.413 0.513 0.411 0.446 0.449 0.097

Table 1 Results from the peak at 522nm. Concentrations in molL-1

Absorption(A) conc. CoCl4

Co(H2O)6 Cl- (Cl-)4 Keq ΔG (J)

Temp(C)

Kelvin Blank 1 2 3 Avg. Avg - blank

abs/ε.l

24 297.15 0 0.699

0.714

0.701

0.705

0.705 0.00122 0.0888 0.355

0.0159 0.865 359

30 303.15 0.0160 0.889

0.825

0.855

0.856

0.840 0.00146 0.0885 0.354

0.0157 1.04 -111

41 314.15 0.027 1.38 1.14

1.25 1.26 1.23 0.00213 0.0879 0.351

0.0153 1.59 -1208

51 324.15 -0.004 1.407

1.525

1.377

1.44 1.44 0.00250 0.0875 0.350

0.0150 1.90 -1729

Table 2 Results from the peak at 690nm. Concentrations in molL-1

295 300 305 310 315 320 325 330

-2

-1.5

-1

-0.5

0

0.5

f(x) = − 0.0799283592116012 x + 24.0774002052171

Gibbs Free Energy

Temperature (K)

G(k

J/m

ol)

Fig 1. Graph shows change in Gibbs free energy with temperature, where slope=- SΔ and the intercept= H.Δ

DiscussionThe experiment had to be restarted as inconsistency in pipetting techniques caused a visible difference in the reaction. The placement of the cuvettes correctly in the spectrometer was very important as putting one in the wrong direction caused a notably large increase in the absorption of the experimental blank. However, being able to pick this out on the graph on the computer allowed the blank to be re-run correctly without affecting the results.Assumptions made in the experiment:

The peaks solely represented CoCl4 and [Co(H2O)6]. No contamination from other transition metal complexes. On small temperature scales S and H do not change significantly but G does.Δ Δ Δ

A discussion can be made using the valid results from the booklet.The reaction moves to produce more CoCl4 at higher temperatures so from the equation G= H-T S the entropy is more of a driving factor for spontaneity.Δ Δ ΔUsing the Beer-Lambert equations, concentrations of CoCl4 were calculated:Absorption= .c.L (1)ε The total cobalt concentration from what was dissolved was 0.09M so Co(H2O)6

concentration could then be calculated. Since Cl- was in excess from the HCl added, the equation for the reaction was used (equation 2 from introduction):Co(H2O)6+4Cl-= CoCl4 +6H2O Thus the concentration of Cl- at equilibrium was 4 times the concentration of Co(H2O)6. The equilibrium constant Keq was then calculated using equation 3:Keq=[CoCl4]/[Co(H2O)6]*[Cl-]4 (3) and then the gibbs free energy change from equation 4:ΔG= RTln(Keq) (4)

A graph of gibbs energy against temperature was plotted and using the equation for Gibbs free energy, equation 5, the thermodynamic constants ΔH and ΔS were calculated:ΔG= ΔH-TΔS (5)Enthalpy=22.4kJ/molEntropy=79.9J/mol/KThe reaction is endothermic with an increase in disorder(entropy).

A question from the booklet discussed the possibility of using a water bath to make a G of 25.0 kJ/mol. Using the results provided in the booklet and Δextrapolating the graph derived from these, y = -0.0799x + 22.4 (where x= temperature and y= G)ΔIf y= 25, x=-32.5This temperature in Kelvin is negative and therefore not possible.

ConclusionThe techniques learnt were carried out effectively to give the correct predicted spectra. The results showed with increasing temperature, the Gibbs free energy of the reaction producing CoCl4 from [Co(H2O)6] decreased. Assuming that enthalpy and entropy remain constant over small temperature ranges these were derived to be 22.4kJ/mol and 79.9J/mol/K respectively.

ReferencesDeGrand, M.J., Abrams, M.L., Jenkins, J.L., Welch, L.E. (2011) Gibbs Energy Changes during Cobalt Complexation: A Thermodynamics Experiment for the General Chemistry Laboratory, Journal of Chemical Education (634-636).

Determination of Ksp, G, H and S for the dissolution ofΔ Δ Δ calcium hydroxide in water

Robert Yates, Prashant Mahendran, Mathew KirbySubmitted 13/05/2015

Edited Version: Prashant MahendranAbstractThe titration experiment allowed the calculation of the apparent equilibrium constants for solubility products of the dissolution of calcium hydroxide in water at 24C and 91C with values of 1.49 x 10-5 and 3.80 x 10-6 respectively. This enabled the calculation of the Gibbs Free Energy, enthalpy and entropy changes of the reaction. The entropy and enthalpy did not vary significantly over the temperature change and were -153 J/mol/K and -18 kJ/mol respectively. The Gibbs Energy was calculated as 27.5 kJ/mol at 24C and 38.0 kJ/mol at 93C. IntroductionThe aim of this experiment is to determine the solubility of calcium hydroxide in water at two temperatures, room temperature (24C) and 93C. A third attempt was completed at 67C. The calcium hydroxide dissolution is shown in equation 1.(1 )Ca¿Where Ca(OH)2 is calcium hydroxide, Ca2+ is calcium ion and OH- is the hydroxyl ion.The apparent equilibrium constant (Ksp) for equation (1) is found from the mass action expression shown in equation 2.(2) K sp=¿ Where [X] represents the concentration of constituent X.To calculate the concentrations for Ksp, titration experiment was undertaken at the three temperatures discussed above, where 0.1M HCl solution was added to a filtered solution of dissolved Calcium Hydroxide until all the hydroxyl had been countered by a proton (H+) from the dissociated HCl. This was measured through using bromothymol blue indicator. This indicator changes colour from blue to yellow once the H+ had countered the OH-. As one H+ is required to counter one OH- the [H+] = [OH-], and as there is one Ca2+ ion per two OH- ions, then [Ca2+] =[OH-]/2. The molar solubility can also be calculated through [OH-]/2.The concentration of [OH] was calculated by using the volume and concentration of the HCl and volume of dissolved Ca(OH)2 as shown in equation 3. The concentration of HCl could be used as the H+ concentration, it is a strong acid, and therefore completely dissociates. (3) mOH- = (mHCl*VHCl)/VOH-

Where mOH- and mHCl is moles of OH- and moles of HCL respectively and VHCl and VOH- represents the volume of the HCl and OH- respectively

The thermodynamic properties of the reaction could then be calculated. This is through the use of Gibbs Free Energy equation shown in equation 4 which allows for the calculation of the Gibbs free energy (∆G) and by rearranging equation 5, and solving for lnKsp with measurements taken at two temperatures, as shown in equation 6, we can determine the enthalpy and entropy.

(4) ∆G=−RTln K sp

Where ∆G is Gibbs Free Energy, R is the ideal gas constant (9.314J/mol/K) and lnKsp represents the natural log of the apparent equilibrium constant for the solubility products.(5) ∆G=∆ H−T ∆ SWhere ∆ H is the enthalphy, T is the absolute temperature in Kelvin and ∆ S is the entropy(6) H = -Rln(KΔ sp(T2)/Ksp(T1)/(1/T2-1/T1)Where T1 = 297.15 Kelvin (24C) and T2 = 366.15Kelvin (93C)Ksp(T1) and Ksp(T2) represents the equilibrium constant for solubility products at temperatures T1 and T2.Experimental Aims:The aims of the experiment included learning that there are two different equilibrium constants, one is apparent where the concentration is being used, and one is thermodynamic, where the activities of the constituent is being used. The activities take into account that the constituent may interact with other species in the solution and not just the one of interest, and therefore tee thermodynamic equilibrium constant is much more accurate than the apparent equilibrium constant.Another aim was to learn about quality control. This was taken into account by doing the titration three times to provide an average, and give us an idea about the precision of the results (how close the results match each other). When looking at the 90C experiment, the filtering of the three samples had to be done at the same time to make sure that the filtering occurred at the same temperature, otherwise the results would be skewed. This required teamwork.

Method1) A 0.1M HCl solution was prepared (already provided) and poured into

burette.2) A solution of Ca(OH)2 at room temperature was stirred3) 10ml of the solution was removed using a syringe and filtered into a

125ml Erlenmeyer flask using filter paper4) 25ml of distilled water and a few drops of bromothymol blue indicator to

the 125ml Erlenmeyer flask5) The 0.1M HCl solution was dripped slowly into the Erlenmeyer flask until

the colour changed from blue to yellow, the volume of HCl was then recorded

6) Steps 3-5 were repeated twice measuring the temperature on the last repeat

7) 100ml of distilled water was then brought to the boil.8) 2g of the Ca(OH)2 was added and left for 5 minutes just below boiling

temperature, and stirred occasionally.9) Steps 3 to 5 were then repeated three times, with the exception that the

samples of Ca(OH)2 had to be filtered at the time so that it was done at the same temperature.

10)The equations discussed in the introduction section were then used to calculate the concentration of the Hydroxyl and calcium ions, and the Gibbs energy, enthalpy and entropy.

Results

Temperature (K)

Average HCl Added (ml)

Hydroxide Ion Concentration

(mol/L)

Calcium Ion Concentration

(mol/L)

Molar Solubility of

Ca(OH)2 Ksp297.15 3.1 3.10E-02 1.55E-02 1.55E-02 1.49E-05340.15 2.03 2.03E-02 1.02E-02 1.02E-02 4.20E-06366.15 1.97 1.97E-02 9.83E-03 9.83E-03 3.80E-06

Table 1 shows the equilibrium constants for the concentrations calculated from the titration data at 3 different temperatures (24°C, 67°C and 91°C)

Temperature (K) G(KJ/mol)Δ H (kJ/mol)Δ S (J/mol/K)Δ

297.15 27.5 -18 -153

340.15 35.0 -18 -156

366.15 38.0 -18 -153

Table 2 shows the Gibbs free energy ( G), enthalpy ( H) and entropy ( S) Δ Δ Δcalculated from the experiment.

Temperature (K) ΔG(KJ/mol) ΔH (kJ/mol)

298.15 33.0323.15 37.4348.15 42.3373.15 47.9

Table 3 displays results from Yeatts and Marshall (1967)

250 270 290 310 330 350 370 3900

5

10

15

20

25

30

35

40

f(x) = 0.154940755855769 x − 18.3390546367114

Gibbs Energy

Temperature (Kelvin)

Gib

bs E

nerg

y

Fig 1 shows the experimentally derived Gibbs free energy against the absolute temperature (K). The intercept represents the enthalpy in (kJ/mol) whilst the gradient represents the entropy (in kJ/mol/K).

DiscussionThe G values calculated for the experimental temperatures were always Δpositive. Therefore the dissolution reaction of Ca(OH)2 in water is not spontaneous for the temperatures used (does not dissolve readily).The Ksp found experimentally was 1.49 x 10-5. The Ksp at room temperature (Yeatts and Marshall 1967) was determined to be 9.37 × 10-6. The error is 37.1%. This error can be attributed to the fact that concentrations were used in this experiment to find the apparent solubilities whereas in the paper, activities were used to find a more accurate result. Furthermore, more time should have been given to allow the calcium hydroxide to equilibrate with the water and dissolve completely. The results show a decrease in solubility with an increase temperature. This is contrary to most solids, however it confirms results found in previous papers (Yeatts and Marshall 1967). The solution was filtered before titrating as any residual undissolved Ca(OH)2 would have continued to react with the HCl, delivering a result that indicated a higher OH- concentration than that was initially dissolved at the initially set temperature. This would have increased the apparent equilibrium constant. Ca(OH)2+ CO2= CaCO3 + H2O The solubility of CO2 decreases as temperature increases (Dodds et al. 1956). As a result, by boiling the water initially, CO2 is driven off from the solution and therefore less CaCO3 is present, which is 100x less soluble in cold water than Ca(OH)2, therefore preventing a decrease in determined solubility.The entropies and enthalpies shown in the table were calculated numerically from equations in the student booklet. The apparent discrepancy from the results at 67 degrees Celsius (the middle experiment) is due to insufficient time given for equilibrium of the dissolution reaction at that temperature. Using graphically determined enthalpy and entropy of -18kJ/mol and -155J/mol/K respectively, the molar solubility of Ca(OH)2 at 50 degrees Celsius was determined to be 0.0123M. These graphically determined values show a

decrease in enthalpy (dissolution of Calcium Hydroxide is exothermic) and a decrease in entropy (dissolution of Calcium Hydroxide results in a decrease in disorder). The literature (Yeatts and Marshall 1967) values for enthalpy and entropy vary over a temperature range of 25-100˚C. The entropy values range from -139 to -208 J/mol/K and the enthalpies range from -12.7 to -35.8 kJ/mol with an increase in temperature. The direction of the entropy and enthalpy changes found in the experiment match the literature and the magnitudes are in the same order. However the variation in the literature values can be attributed to the increase in order of the water molecules around the Ca2+ and the OH- ions (Yeatts and Marshall 1967). This links to the fact that the experiment was carried out using concentrations rather than activities. The enthalpy change for the dissolution reaction is negative, so the reaction is exothermic, so due to Le Chatelier’s principle, a decrease in temperature causes the equilibrium to shift to dissolve more calcium hydroxide to increase the temperature of the system. This causes the trend in decreasing solubility with increasing temperature for Calcium Hydroxide. Hydroxide Ksp at 25 degrees CelsiusMg(OH)2 5.6 x 10-12Ca(OH)2 4.7 x 10-6Sr(OH)2 6.4 x 10-3Table 4 shows a trend with an increase in solubility product from Magnesium to Strontium, so the solubility increases further down the Group in the periodic table (Group 2). The atomic radius and amount of electrons increases down the group so that the metal cation has more electron shielding and so the hydroxyl ion is less strongly bonded to it and therefore more likely to dissociate. As a result, solubility increases down the group. ConclusionThis experiment has shown that the concentrations of hydroxyl in the solution decrease with increasing temperature, from 3.10E-02M at 297.15K to 1.97E-02M at 336.15K. This corresponds to a similar concentration decrease of 1.55E-02M at 297.15K and 9.83E-03M at 336.15K for the calcium ion. This is due to a decrease in solubility with temperature for calcium hydroxide and thus an increase in Gibbs free energy (the reaction is less spontaneous). At 25 degrees Celsius, the experimentally derives thermodynamic constants are as follows; ΔG of 27.5 kJ/mol, ΔH of -18 kJ/mol and ΔS of -153 J/mol/K. At 93 degrees these were 38kJ/mol, -18 kJ/mol and -153 J/mol/K respectively.

ReferencesDodds, W. S., Stutzman, L. F. & Sollami, B. J. (1956) Carbon Dioxide Solubility in Water Journal of Chemical and Engineering Data, 1 92-95.Yeatts, L.B., Marshall (1967), W.L. Aqueous Systems at High Temperature. XVIII. Activity Coefficient Behaviour of Calcium Hydroxide in Aqueous Sodium Nitrate to the Critical Temperature of Water, J. Phys Chem., 71(8), 2641-2650.

Oxidation Kinetics of Fe 2+ in aqueous solutions Prashant Mahendran, Robert Yates, Matthew Kirby

First version submitted 15/5/15Edited Version: Prashant Mahendran

AbstractOxidisation of Fe(II) to Fe(III) in aqueous solutions have been analysed through the use of UV/Visible light spectroscopy. The experiment results gave the changes in Fe(II) concentrations over time, allowing calculation of apparent reaction rates to produce results showing an increase in the reaction rate with pH with apparent reaction rate constants of 0.0418s-1 at pH 7.17 to 0.219s-1 at pH 7.55.IntroductionIron in aquatic systems can exist in several oxidation states: metallic iron (iron metal), ferrous iron (Fe2+) and ferric iron (Fe3+). This oxidation state depends on the amount of dissolved oxygen in aquatic systems. The importance of iron oxidation within an aquatic system is great due to the fact that Fe (III) forms hydroxides. This is insoluble and when in an aquatic environment can suffocate fish due to it getting lodged in gills and thus suffocating them. The experiment involved analysing the iron concentrations by using both a colour changing which starts off purple in the presence of Fe(II) and decreases to colourless with no Fe(II), as well as using UV/Vis to calculate the absorbance of light and therefore the concentration. The more of the Fe(II) present, the more light was absorbed, and therefore the higher the concentration Fe(II).The first step of the experiment involved creating a calibration curve using 1ml samples with known variations in iron concentration, by using a 0.02M of iron Fe(NH4)2(SO4)2.6H2O and diluting it with distilled water, with 0.2ml steps (so 0ml water 1ml 0.02M Fe through to 1ml of water and 0ml Fe).

The experiment used 490ml of water kept at a stable ph. 1ml of the iron solution was added to this beaker and the beaker was oxygenated. This allows us to determine how concentrations of Fe(II) varied over time due to oxidation to Fe(III) by taking samples at different times. The concentration was measured using the UV/Visible light spectrometry. This was used to measure the absorbance of the iron solution and therefore the change in concentration of the Fe(II) over time. The samples should show that Fe(II) concentration decreases with time as the iron is oxidised.The rate for the oxidation of Fe(II) is shown in equation 1:r= -d[Fe(II)]/dt= k[Fe(II)][OH-]2pO2 (1)Where k is the true rate constant. When pH (and therefore OH- concentration) and pO2 are kept constant as done in the experiment, the equation reduces to equation 2:r= -d[Fe(II)]/dt= k1[Fe(II)] (2)Where k1 is the apparent rate constant for the pseudo first order rate reaction. In the experiment, the natural log of the concentrations of Fe(II) will be plotted against time to show a linear relationship as shown in this rearrangement of equation 2:ln([Fe(II)])=-k1t +c (3)Where t is time and k and c are unknown constants.

The gradient of the graph corresponds to k1 and thus k1 can be derived experimentally. Objectives

1. An Fe(II) solution will be oxidised to Fe(III)2. 5ml of the reaction mixture will be pipetted into a test tube, forming a purple

complex with ferrozine and halting the reaction.3. This will be repeated at regular time intervals.4. Using UV/Vis spectrophotometry at 1 wavelength (562nm), the absorbance will

be used as an indication of Fe(II) content which decreases over time.5. This is used to find the rate constant of the reaction and repeated at 2 pH values.

Method1. Prepared 5 cuvettes of solutions 1mL with differing concentrations (1.0, 0.8, 0.6, 0.4 and 0.2mL) of ferrozine (Fe2+) to H20. We also used a blank H20 solution to calibrate the instrument before analysis.2. Analysed each cuvette to plot an initial calibration curve of absorbance at 562nm.3. Prepared 10 X 10 ml test tubes by adding 1 ml of acetate buffer and 0.1ml of ferrozine solution. Labelled the solutions from t0 to t9 (t = time). Also prepared was a solution labelled blank.4. Add 490 ml of DD water to 2.5 ml of 1M imidazole buffer and used a magnetic stirrer to ensure the solution is well mixed.5. The solution is then purged with air and oxygen for at least 5 minutes before the experiment begins. A pH electrode is calibrated and put into the reaction vessel to measure the pH.6. A chosen amount (1.5ml) of NaOH is added to raise the pH to the desired 7.2.7. Prepare a 5 ml pipette for the first sample collection at t0.8. Once the pH is at 7.2 the reaction is initiated by adding 1 ml of the Fe(II) solution, take immediately the first sample at t0.9. Note the pH when each 5 ml sample is pipetted out and note any colour change. We can use the colour change of the test tubes and of the solution to determine the time interval of pipetting.10. Once all samples have been collected the absorption measurements can be taken using the UV/V spectrophotometer.11. Repeat steps 3-10 using an initial solution pH of 7.6.

ResultsOriginal [Fe2+] Vol Fe

solutionVol H2O [Fe2+] Absorbance at

562nm0.001 0 1 0 1.00E-03

0.001 0.2 0.8 0.0002 5.21E-01

0.001 0.4 0.6 0.0004 1.09E+00

0.001 0.6 0.4 0.0006 1.66E+00

0.001 0.8 0.2 0.0008 2.19E+00

0.001 1 0 0.001 2.67E+00

Table 1 Data for calibration curve

Fig 1 Calibration curve for [Fe(II)] concentrations, showing error bars of one standard deviation

pH Time (mins) Abs. [Fe2+] ln[Fe2+]7.2 t0 0 0.34 1.36E-04 -8.90E+00

7.17 t1 1 0.355 1.42E-04 -8.86E+007.18 t2 2 0.353 1.41E-04 -8.86E+007.17 t3 5 0.33 1.32E-04 -8.93E+007.17 t4 7 0.317 1.27E-04 -8.97E+007.17 t5 11 0.283 1.13E-04 -9.08E+007.17 t6 14 0.241 9.66E-05 -9.24E+007.15 t7 19 0.184 7.38E-05 -9.51E+007.16 t8 21 0.17 6.82E-05 -9.59E+007.15 t9 23 0.136 5.46E-05 -9.82E+00

Table 2 change in [Fe(II)] concentration with time at average pH 7.17.

pH Time (mins) Abs. [Fe2+] ln[Fe2+]7.57 t0 0 0.346 1.39E-04 -8.88E+007.55 t1 0.5 0.322 1.29E-04 -8.96E+007.56 t2 1 0.316 1.27E-04 -8.97E+007.55 t3 2 0.301 1.21E-04 -9.02E+007.54 t4 4 0.21 8.42E-05 -9.38E+007.54 t5 5 0.141 5.66E-05 -9.78E+007.54 t6 6 0.122 4.90E-05 -9.92E+007.54 t7 7 0.087 3.50E-05 -1.03E+017.54 t8 8 0.062 2.50E-05 -1.06E+017.54 t9 9 0.049 1.98E-05 -1.08E+01

Table 3 change in [Fe(II)] concentration with time at average pH 7.55.

0 1 2 3 4 5 6 7 8 9 10

-1.20E+01

-1.00E+01

-8.00E+00

-6.00E+00

-4.00E+00

-2.00E+00

0.00E+00

f(x) = − 0.218986903890913 x − 8.73023158505323R² = 0.96726724849411

f(x) = NaN x + NaNR² = 0 Concentration of Fe(II) against time

pH 7.17Linear (pH 7.17)pH 7.55Linear (pH 7.55)

Time (mins)

ln F

e(II

)

0.00 0.50 1.00 1.50 2.00 2.50 3.000

0.0002

0.0004

0.0006

0.0008

0.001

0.0012

f(x) = 0.000369100571580399 x − 1.92791253369995E-07R² = 0.99936616188047

Calibration Curve

Absorbance (A)

[Fe2

+]

con

cen

trat

ion

(M

)

Fig 2 Graph of natural log of [Fe(II)] against time (mins). The anomalous first data point for pH 7.17 has been excluded from the graph due to insufficient mixing time.

Fig 3 Graph of apparent rate constant k1 against pH. k1=-m where m is the gradient of each of the trendlines in Fig 2.DiscussionThe experiment based the derivation of Fe(II) concentrations on a calibration curve with a linearity of 0.9994 which is very close to 1 and therefore accurate. The standard deviation for the Fe(II) concentrations determined by the calibration curve is 5.67 x 10-5 which is and order of magnitude smaller than the results obtained. This is relatively small however preferably in future a calibration curve with more measurements at smaller absorbance values would be made to improve accuracy of the determined concentrations. The experimental blank was 0.002A. Using equation 4:ydl=yblank+3s (4)Where ydl is the minimum detectable signal and s is the standard deviation, the ydl for this calibration curve is 0.00117, which corresponds to a minimum detectable concentration of 4.68x10-7. None of the results fall close to this value, so we can accept the results as being above the minimum detection limit. Fig 2 shows how at a higher pH causes an increase in the rate of oxidation of Fe(II) to Fe(III). This is because of a higher hydroxyl ion concentration at a higher pH (see

0 1 2 3 4 5 6 7 8 9 10

-1.20E+01

-1.00E+01

-8.00E+00

-6.00E+00

-4.00E+00

-2.00E+00

0.00E+00

f(x) = − 0.218986903890913 x − 8.73023158505323R² = 0.96726724849411

f(x) = NaN x + NaNR² = 0 Concentration of Fe(II) against time

pH 7.17Linear (pH 7.17)pH 7.55Linear (pH 7.55)

Time (mins)

ln F

e(II

)

7.17 7.550.00E+00

5.00E-02

1.00E-01

1.50E-01

2.00E-01

2.50E-01

k1 vs pH

pH

k1

(/s

)

equation 3). The pH difference of 0.38 units causes an increase in the apparent rate constant (k1) of 0.0418s-1 at pH 7.17 to 0.219s-1 at pH 7.55. The O2 drives the oxidation reaction, while increasing pH means that there are more hydroxyl groups present, and because the Fe(II) will be oxidised to Fe(III) and form the iron hydroxide Fe(III)(OH)3.Fe(II) content in natural waters depends on the biological and chemical conditions which control the variables in the rate equation. For example, groundwater without presence of dissolved oxygen could undergo anoxic reduction of iron from Fe(III) to Fe(II) mediated by bacteria. While in a river, the surface agitation and contact with air would suggest more oxidising conditions and a much higher presence of Fe(III), depending on the level of biological activity. The trend shown in Emmenegger et al. (1998) of apparent rate constants in a lake varying with pH shows the same trend as our results (increasing rate constant with increasing pH) however our values, 0.0418s-1 at pH 7.17 to 0.219s-1

at pH 7.55 are lower in comparison to the results from the lake experiment, which are approximately 0.287s-1 at pH 7.17 and 0.407s-1 at pH 7.55. This could be due to acceleration of the oxidation of Fe(II) by an organic ligand in natural systems (Emmenegger et al. 1998).

ConclusionAnalysis of results gained from the experiment have shown how the oxidation of Fe(II) to Fe(III) increase with pH and therefore concentration of [OH]- while the presence of dissolved oxygen drives the reaction. The experiment has provided apparent reaction rate constants of 0.0418s-1 at pH 7.17 to 0.219s-1 at pH 7.55

ReferencesClarke, W.A. et al. (1997). Ferric hydroxide and ferric hydroxysulfate precipitation by bacteria in an acid mine drainage lagoon, FEMS Microbiology Reviews, 20, 351–361Johnson, D.B., and Hallberg, K.B. (2005) Acid mine drainage remediation options: a review, Science of the Total Environment, 338, 3–14Emmenegger, L., et al. (1998). Oxidation kinetics of Fe(II) in a eutrophic Swiss lake, Environ. Sci. Technol., 32, 2990-2996.

Oxidation Kinetics of Fe 2+ in aqueous solutions of varying ion content

Independent Lab ReportPrashant Mahendran

AbstractOxidisation of Fe(II) to Fe(III) in aqueous solutions have been analysed through the use of UV/Visible light spectroscopy. This was repeated at 3 pHs for 7g/L NaCl brine solution, Highland Spring water and Evian water, as well as the pure milli-Q water. The apparent rate constant of the oxidation of Fe(II) to Fe(III), k1, has been shown to increase with pH for all solutions. The presence of high concentrations of Cl- reduces the rate of reaction and high alkalinity in the form of HCO3

- (in Evian) appears to cause an increase in the rate of reaction. At pH ~ 7.6 the apparent rate constant k1 is 0.199±0.00904 in Highland Spring Water, 0.278±0.00696 in Evian water, 0.604±0.265 in the brine solution and 0.219±0.0142 in the pure Milli-Q water. IntroductionThis experiment aims to explore the effect of varying the ion content of aqueous solutions on the rate of oxidation of Fe(II) to Fe(III). This follows on from the iron oxidation experiment carried out previously, but 7g/L NaCl brine solution, Highland Spring water and Evian water were used instead of just milli-Q water. This was repeated at 3 pHs for each solution to observe how the rate of the reaction varies with pH in different solutions. The rate for the oxidation of Fe(II) is shown in equation 1:r= -d[Fe(II)]/dt= k[Fe(II)][OH-]2pO2 (1)Where k is the true rate constant. When pH (and therefore OH- concentration) and pO2 are kept constant as done in the experiment, the equation reduces to equation 2:r= -d[Fe(II)]/dt= k1[Fe(II)] (2)Where k1 is the apparent rate constant for the pseudo first order rate reaction. In the experiment, the natural log of the concentrations of Fe(II) will be plotted against time to show a linear relationship as shown in this rearrangement of equation 2:ln([Fe(II)])=-k1t +c (3)Where t is time and k and c are unknown constants. The gradient of the graph corresponds to k1 and thus k1 can be derived experimentally. Ionic strengths, I, of each of the solutions could be calculated using the following equation:I= 0.5Σcz2 (4)Where c is the concentration of each species and z is the charge of that species. The concentrations provided on the labels for Evian and Highland Spring waters could be used (after converting mg/L to mol/L).

IonConcentration mg/L

Evian Highland Spring Brine Milli-QCa 80 40.5 0 0

Mg 26 10.1 0 0Na 6.5 5.6 7000 0K 1 0.7 0 0HCO3 360 150 0 0SO4 12.6 5.3 0 0Cl 6.8 6.1 7000 0HNO3 3.7 3.1 0 0Table 1 Concentrations of ions in the solutions usedPure Milli-Q water has an ionic strength of 0, the brine solution has an ionic strength of 0.120M, the Highland Spring water has an ionic strength of 0.00443M and the Evian water has an ionic strength of 0.00963M.At higher pH Fe(II) is more easily oxidised as it exists as FeOH+ and Fe(OH)2

0 which are more easily oxidised due to the presence of hydroxyl groups (Morgan and Lahav, 2007).The presence of Cl- lowers the oxidation rate by forming complexes with Fe(II) that are less easily oxidised (Morgan and Lahav, 2007). Solutions of high alkalinity will result in the oxidation reaction depending on the solubility of siderite, FeCO3 (Morgan and Lahav, 2007), which will lower the rate of reaction due to carbonate anions forming compelxes which are less easily oxidised (Morgan and Lahav, 2007).

Hypothesis: Evian Water has a higher bicarbonate content so this will cause a decrease in the rate of reaction, and brine contains a high chloride ion content so this will cause a decrease in the rate of reaction. Highland Spring water will be in between. The rate of reaction will increase with pH.

Method1. Prepared 5 cuvettes of solutions 1mL with differing concentrations (1.0, 0.8, 0.6, 0.4 and 0.2mL) of ferrozine (Fe2+) to H20. We also used a blank H20 solution to calibrate the instrument before analysis.2. Analysed each cuvette to plot an initial calibration curve of absorbance at 562nm.3. Prepared 10 X 10 ml test tubes by adding 1 ml of acetate buffer and 0.1ml of ferrozine solution. Labelled the solutions from t0 to t9 (t = time). Also prepared was a solution labelled blank.4. Add 490 ml of DD water to 2.5 ml of 1M imidazole buffer and used a magnetic stirrer to ensure the solution is well mixed.5. The solution is then purged with air and oxygen for at least 5 minutes before the experiment begins. A pH electrode is calibrated and put into the reaction vessel to measure the pH.6. A chosen amount (1.5ml) of NaOH is added to raise the pH to the desired 7.2.7. Prepare a 5 ml pipette for the first sample collection at t0.8. Once the pH is at 7.2 the reaction is initiated by adding 1 ml of the Fe(II) solution, take immediately the first sample at t0.9. Note the pH when each 5 ml sample is pipetted out and note any colour change. We can use the colour change of the test tubes and of the solution to determine the time interval of pipetting.

10. Once all samples have been collected the absorption measurements can be taken using the UV/Vis spectrophotometer, each time checking that the blank solution results in an absorbance very close to 0 (<0.01A)11. Repeat steps 3-10 using an initial solution pH of 7.6 and 8.0. 12. Repeat steps 3-11 with Highland Spring water, Evian water and 7g/L NaCl brine solution.

ResultsOriginal [Fe2+] (M) Vol Fe

(ml)Vol H2O

(ml)[Fe2+] (M) Absorbance at

562nm (A)0.0001 0 1 0 0.0000.0001 0.2 0.8 0.00002 0.1810.0001 0.4 0.6 0.00004 0.3510.0001 0.6 0.4 0.00006 0.5570.0001 0.8 0.2 0.00008 0.7760.0001 1 0 0.0001 0.944

Table 2 Data for calibration curve

0.00 0.10 0.20 0.30 0.40 0.50 0.60 0.70 0.80 0.90 1.000

0.00002

0.00004

0.00006

0.00008

0.0001

0.00012

f(x) = 0.00010413646723772 x + 1.24677725487417E-06R² = 0.998371188046197

Calibration Curve

Absorbance (A)

[Fe2

+]

con

cen

trat

ion

(M

)

Fig 1 Calibration curve for [Fe(II)] concentrations, showing error bars of one standard deviation. The standard deviation for y sy=1.69x10-6, for the intercept

sb=1.20x10-6 and for the gradient sm=2.10x10-6.

Pure Milli-Q water

0 2.5 5 7.5 10

-16.00

-14.00

-12.00

-10.00

-8.00

-6.00

-4.00

-2.00

0.00

f(x) = − 1.9253896060735 x − 10.0016480051753R² = 0.948275951017841

f(x) = − 0.218986903890913 x − 8.73023158505322R² = 0.96726724849411

f(x) = NaN x + NaNR² = 0 Concentration of Fe(II) against time for Milli-Q water

pH 7.17

Linear (pH 7.17)

pH 7.55

Linear (pH 7.55)

pH 8.00

Linear (pH 8.00)

Time (mins)

ln F

e(II

)

Fig 2 Graph of natural log of [Fe(II)] against time (mins) for pure Milli-Q water. The data for pH 7.17 and 7.55 was from the experiment carried out previously, and the data for pH 8.00 was obtained from this experiment using the new calibration curve. 7g/L brine solution

Time (mins) Absorbance (A) [Fe2+] (M) ln[Fe2+]0 0.084* 0.0000094 -11.62 0.308 0.0000318 -10.44 0.259* 0.0000269 -10.55 0.285 0.0000295 -10.47 0.188 0.0000198 -10.8

10 0.086 0.0000096 -11.612 0.067 0.0000077 -11.8

Table 3 Data for brine at average pH 7.18. Asterisked absorbances do not fit the trend; the value at time=0 is anomalous due to insufficient mixing.

Time (mins) Abs. (A) [Fe2+] (M) ln[Fe2+]

0 0.075* 0.0000085 -11.7

1 0.194 0.0000204 -10.8

2 0.055 0.0000065 -11.9

3 0.051 0.0000061 -12.0


Time (mins) Abs. (A) [Fe2+] (M) ln[Fe2+]0 0.038* 0.0000048 -12.2

0.3333 0.143 0.0000153 -11.10.66667 0.098 0.0000108 -11.4

1 0.08 0.000009 -11.61.33333 0.02 0.000003 -12.71.6667 0.161* 0.0000171 -11.0

2 0.097* 0.0000107 -11.4


0 2 4 6 8 10 12 14

-14.0

-12.0

-10.0

-8.0

-6.0

-4.0

-2.0

0.0

f(x) = − 1.52096379736493 x − 10.4472444443919R² = 0.865737843069533

f(x) = − 0.60362306483545 x − 10.3763891453164R² = 0.789342675246269

f(x) = − 0.157662825151846 x − 9.85383383323781R² = 0.93318850535694

Concentration of Fe(II) against time for brine

pH 7.18Linear (pH 7.18)pH 7.58Linear (pH 7.58)pH 8.01Linear (pH 8.01)

Time (mins)

ln[F

e2+

]

Fig 3 Graph of natural log of [Fe(II)] against time (mins) for 7g/L brine solution. Asterisked anomalous absorbances have been excluded from the graph.

Evian Mineral WaterTime (mins) Abs. (A) [Fe2+] (M) ln[Fe2+]

0 0.438 0.0000448 -10.02 0.301 0.0000311 -10.44 0.273 0.0000283 -10.56 0.219 0.0000229 -10.79 0.18 0.000019 -10.9

12 0.138 0.0000148 -11.115 0.106 0.0000116 -11.419 0.086 0.0000096 -11.6

Table 6 Data for Evian at average pH 7.36.

Time (mins) Abs. (A) [Fe2+] (M) ln[Fe2+]0 0.296 0.0000306 -10.4

1 0.247 0.0000257 -10.63 0.123 0.0000133 -11.25 0.061 0.0000071 -11.96 0.048 0.0000058 -12.17 0.033 0.0000043 -12.48 0.025 0.0000035 -12.69 0.017 0.0000027 -12.8



0 0.237 0.0000247 -10.6

0.5 0.123 0.0000133 -11.2

1 0.056 0.0000066 -11.9

1.25 0.039 0.0000049 -12.2

1.5 0.029 0.0000039 -12.5

1.75 0.01 0.000002 -13.1

2 0.009 0.0000019 -13.2

2.25 0.004 0.0000014 -13.5


0 2 4 6 8 10 12 14 16 18 20

-16.0

-14.0

-12.0

-10.0

-8.0

-6.0

-4.0

-2.0

0.0

f(x) = − 1.30734931497898 x − 10.60255399946R² = 0.989558398660631f(x) = − 0.277467491614516 x − 10.378126848607R² = 0.996241021737718f(x) = − 0.078266363659431 x − 10.1518736335243R² = 0.980142105328997

Concentration of Fe(II) against time for Evian Mineral Water


Time (mins)

ln[F

e2+

]

Fig 4 Graph of natural log of [Fe(II)] against time (mins) for Evian Mineral Water.

Highland Spring Mineral Water


0 0.253* 0.0000263 -10.5

3 0.319 0.0000329 -10.3

6 0.311 0.0000321 -10.3

10 0.276 0.0000286 -10.5

14 0.249 0.0000259 -10.6

19 0.231 0.0000241 -10.6

22 0.233 0.0000243 -10.6

25 0.209 0.0000219 -10.7

Table 9 Data for Highland Spring Water at average pH 7.22. Asterisked absorbances do not fit the trend; the value at time=0 is anomalous due to insufficient mixing.

Time (mins) Abs. (A) [Fe2+] (M) ln[Fe2+]0 0.282 0.0000292 -10.41 0.261 0.0000271 -10.5

2.5 0.195 0.0000205 -10.84 0.137 0.0000147 -11.1

5.5 0.118 0.0000128 -11.37 0.07 0.000008 -11.7

8.5 0.05 0.000006 -12.010.5 0.027 0.0000037 -12.5

Table 10 Data for Highland Spring Water at average pH 7.62.


0 0.434 0.0000448 -10.0

30 0.119 0.0000311 -10.4

45 0.077 0.0000283 -10.5

1 0.061 0.0000229 -10.7

1.25 0.06 0.000019 -10.9

1.5 0.05 0.0000148 -11.1

1.75 0.045 0.0000116 -11.4

2.25 0.044 0.0000096 -11.6

Table 11 Data for Highland Spring Water at average pH 7.97.

0 2 4 6 8 10 12

-14.0

-12.0

-10.0

-8.0

-6.0

-4.0

-2.0

0.0

f(x) = − 0.737931034482759 x − 9.9948275862069R² = 0.989456399619739f(x) = − 0.199140180453664 x − 10.3308329509311R² = 0.987775130134237

f(x) = NaN x + NaNR² = 0 Concentration of Fe(II) against time for

Highland Spring Mineral water


Time (mins)

ln[F

e2+

]

Fig 5 Graph of natural log of [Fe(II)] against time (mins) for Highland Spring Water. The asterisked anomalous absorbance has been excluded from the graph.

Rate constant k1

pH k1Highland

7.22 0.0182±0.003867.62 0.199±0.009047.97 0.738±0.0311

Evian7.36 0.0783±0.004557.61 0.278±0.006967.95 1.31±0.0548

brine7.18 0.158±0.05637.58 0.604±0.2658.01 1.52±0.385

Milli-Q7.17 0.0418±0.003527.55 0.219±0.0142

8 1.93±0.225Table 12 rate constants k1 for different solutions at varying pH. k1=-m where m is the gradient of each of the trendlines in the previous graphs, shown with ±standard deviations in gradient of the graphs.

7.1 7.2 7.3 7.4 7.5 7.6 7.7 7.8 7.9 8 8.10

0.5

1

1.5

2

2.5

Rate constant k1 against pH

Highland SpringEvianBrineMilli-Q

pH

k1

Fig 6 Graph of apparent rate constant k1 against pH. k1=-m where m is the gradient of each of the trendlines in the previous graphs.

Fig 7 Test tubes from experiment showing trend of lightening of purple colour from t0 to t7 indicating decreasing Fe(II) over time as it oxidises to Fe(III).

DiscussionQuality control: The errors in y, intercept and gradient for the calibration curve are so small they were negligible and not propagated through the rest of the experiment. Standard deviations in gradient of the graphs used to determine rate constants are included in table 10 and considered when comparing values. Repeats were carried out approximately one every 10 cuvettes in the UV/Vis spectrophotometer by pouring more solution out of the test tube into a cuvette again and running it in the machine again. An average of the consistent values obtained was then used and displayed in the tables above. Experimental blank solutions containing all the constituents of the other test tubes except the purple ferrozine complex were used and run in the machine, ensuring blank readings close to zero before proceeding with the experiment each time a set of test tubes were brought to the spectrophotometer. Due to the well buffered nature of Evian Water (high bicarbonate content) the pH crept up significantly during the experiment and eventually the experiments had to be repeated whilst adding imidazole buffer at regular intervals throughout the experiment to keep the pH stable. Fig 6 shows how at a higher pH causes an increase in the rate of oxidation of Fe(II) to Fe(III). This is because of a higher hydroxyl ion concentration at a higher pH (see equation 1). At higher pH Fe(II) is more easily oxidised as it exists as FeOH+ and Fe(OH)2

0 which are more easily oxidised due to the presence of hydroxyl groups (Morgan and Lahav, 2007).As the pH increases, the pure Milli-Q water’s rate constant increases to eventually be much higher at pH 8 than all other solutions. This is because of the increase in hydroxyl ion concentration at a higher pH increasing the rate of reaction (see equation 1) forming FeOH+ and Fe(OH)2

0 which are more easily oxidised, without being affected by other ions; the concentration of OH- causes an increase in reaction rate in mill-Q water most strongly as the activity of OH- is not significantly reduced by the presence of other ions. Pure Milli-Q water has an ionic strength of 0, the brine solution has an ionic strength of 0.120M, the Highland Spring water has an ionic strength of 0.00443M and the Evian water has an ionic strength of 0.00963M. All the solutions except Milli-Q water exceed the threshold of I=0.002M where the activities of the ions become significantly lower than concentrations.The trend in ionic strength goes from Milli-Q water, Highland Spring water, Evian water and then the brine solution. From Fig 6 and Table 12 it can be seen that the rate of oxidation is consistently higher in Evian water than Highland Spring for every pH(but the low pH value of evian was higher than that of Highland Spring), and the rate of oxidation is consistently higher in brine than Evian water for every pH(except pH 8 where the standard deviations overlap slightly).The presence of Cl- in the brine solution lowers the oxidation rate by forming complexes with Fe(II) that are less easily oxidised, as expected (Morgan and Lahav, 2007). The high bicarbonate content in Evian resulted in the oxidation reaction depending on the solubility of siderite, FeCO3 (Morgan and Lahav, 2007), which was expected to decrease the rate of reaction relative to Highland Spring water however the rate of reaction was higher in Evian water than Highland Spring water which is not fully understood. This may be due to HCO3

- ions accelerating the reaction rather than slowing it down by forming CO3

2-.

It is very clear that the graph of k1 against pH for brine solution does not have the same shape as the other solutions, this is likely due to the insufficient data gained from the experiments for brine solution resulting in initial graphs of ln[Fe(II)] against time with poorly fitting trendlines (low R2 values). However this may also be due to the effect of such a high ionic strength of 0.120M.

ConclusionThe apparent rate constant of the oxidation of Fe(II) to Fe(III), k1, has been shown to increase with pH for all solutions, due to the formation of FeOH+ and Fe(OH)2

0 which are more easily oxidised due to the presence of hydroxyl groups (Morgan and Lahav, 2007). This increasing trend is affected by the presence of high ionic strengths in other solutions as the trend is strongest for pure water. The presence of high concentrations of Cl- reduces the rate of reaction by forming less easily oxidised complexes with Fe2+ ions (Morgan and Lahav, 2007). High alkalinity in the form of HCO3

- (in Evian) appears to cause an increase in the rate of reaction.

ReferencesMorgan, B., Lahav, O. (2007), The effect of pH on the kinetics of spontaneous Fe(II) oxidation by O2 in aqueous solution – basic principles and a simple heuristic description., Chemosphere, 68, 2080-2084.

Overall Conclusion

The lab report has given me a deeper understanding of laboratory procedures as well as the theory of the chemistry behind the range of experiments undertaken. I learned skills such as using a UV/Vis spectrophotometer, pipetting, and accurately performing a titration.

environmental geochemistry course lab report

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