Chemistry 26.1 2nd Written Laboratory Examination Reviewer

A.Y. 2010 – 2011

Experiment # 8 QUANTITATIVE DETERMINATION OF TOTAL HARDNESS IN DRINKING WATER BY

COMPLEXOMETRIC EDTA TITRATION INTRODUCTION Primary Standard: Calcium Carbonate (CaCO3) Titrant: EDTA Titrant/Analyte: Mineral Water Sample Indicator: Eriochrome Black T (EBT) Total hardness is a measure of the concentration of calcium and magnesium ions in water. It is, however, measured primarily as ppm of calcium carbonate (CaCO3) due to the negligible presence of other ions. Ideal Freshwater Ca2+ Concentration: 10 – 250 ppm Ideal Freshwater Mg2+ Concentration: 5 – 125 ppm Choice of Titrant: Ethylenediaminetetraacetic acid (EDTA) EDTA is set as the titrant due to its hexadentate property as a ligand in the formed metal-EDTA complex. This means that it can bind with metal ions in a 1:1 ratio due to its numerous binding sites, simplifying calculations. It is important to note that EDTA exists in different states depending on the pH of the solution. At pH 10, EDTA is represented as Y4- in equations. A 1:1 ratio also means a sharper endpoint for the titration.

Choice of Indicator: Eriochrome Black T (EBT) Eriochrome Black T (EBT) was used in the experiment to determine the endpoint of the EDTA titration process. The color of the Metal-EBT complex in a pH 10 solution (due to ammonia buffering) is wine red*. Upon reaching the endpoint, the complex is broken up and the solution appears blue. Note that the charge of In is -3.

H2Y2- + MgIn- -> MgY2- + HIn2- + H+

One limitation of the EBT indicator, however, is that it readily undergoes air oxidation, so the solution must be covered when not in use. Calmagite is a possible substitute that, unlike EBT, does not undergo air oxidation. *For a pH value less than 6.3, the transition is wine red to red. (Hard to detect) *For a pH value greater than 11.6, the transition is wine red to orange. (Hard to detect) SOLUTIONS AND REAGENTS Stock EDTA Solution Standardization Ratio: 1 mol EDTA: 1 mol Ca2+ (CaCO3 standard) The solubility of EDTA crystals in water is quite limited as witnessed in the experiment. This is why heat must be introduced while stirring to ease dissolution. Note that this is an endothermic system. Magnesium ions (Mg2+) is added to ensure a sharp endpoint during titration. Without them, a premature endpoint may result. Sodium hydroxide pellets are added to dispel turbidity (murkiness) in the solution. Excessive addition may upset the pH balance of the solution and result in the formation of magnesium hydroxide (Mg(OH)2).

wine red blue

NOTE: Calcium hydroxide (Ca(OH)2) will not precipitate because of its relatively high Ksp value, meaning it favors the ionic form. Calcium Carbonate Solution (CaCO3) NH3-NH4+ Buffer The buffer is added to the solution to maintain a pH of 10. Ammonia was chosen because its pKa value is approximately 9.26, very close to the desired value of 10. This is due to the following reasons: 1) Sharp Endpoint 2) Increased Ca2+ and Mg2+ Selectivity 3) Wine Red to Blue Transition Ammonia, due to its volatile nature, may evaporate if the solution is left uncovered. This will result in a decrease in the pH value. Excessive addition of the buffer may lead to a decrease in the sharpness of the endpoint. SAMPLE ANALYSIS As a hexadentate ligand, EDTA forms highly stable complexes with metal ions in the titration process, which is one of the fundamental principles of the experiment. The Kf (Complex Formation Constant) hierarchy is as follows,

Ca-EDTA > Mg-EDTA > MgIn- > CaIn-

More About Endpoint Essentially, what happens at the endpoint is that all the MgIn- complexes are replaced with Mg-EDTA complexes, causing the shift to the blue indicator color. Should a violet color change result, potassium cyanide (KCN) is added after the buffer to act as a masking agent and prevent interference of iron.

HIn2- + Fe3+ -> FeIn- + H+

6CN- + Fe3+ -> [Fe(CN)6]3-

This is possible because Kf of the Fe-CN complex is larger than that of the Fe-In complex. Be warned, however, that although increasing the pH to 12 yields a sharper endpoint, the formation of magnesium hydroxide (Mg(OH)2) is imminent and undesirable, causing much error. OH- may result from the hydrolysis of the ammonia buffer in the solution. CALCULATIONS Molarity of EDTA Titrant

[EDTA] = ( )

Remember that whatever you put for the aliquot factor has to make the numerator smaller than the original since you diluted the solutions. CaCO3 Titer

Titer = (

)

Parts Per Million (PPM) CaCO3

ppm CaCO3 = (

)

violet

Experiment # 9 REDOX TITRATION: WINKLER METHOD FOR DISSOLVED OXYGEN DETERMINATION

INTRODUCTION Primary Standard: Potassium Iodate (KIO3) Titrant: Sodium Thiosulfate (Na2SO3) Titrant/Analyte: Dissolved Oxygen Indicator: 1% Starch Solution Limiting Reactant: Dissolved Oxygen (O2) The Winkler method is a redox titration method employed to test the life sustainability of pond water. It makes use of several reduction-oxidation processes that relate the dissolved oxygen content to the amount of titrant used. In this particular case, manganese (Mn2+) is both reduced and oxidized at certain parts of the experiment. STANDARDIZATION Stoichiometric Ratio: 1 mmol IO3

- = 6 mmols S2O32-

KIO3 is initially placed in the beaker, and water, potassium iodide (KI) and sulfuric acid (H2SO4) are eventually added to it in the order specified, yielding the following products in a dark yellow solution.

IO3- + 8I- + 6H+ -> 3I3- + 3H2O

Sulfuric acid in particular was added due to the need for an acidic environment for the reaction with iodate (IO3

-) to take place. The triiodide formed on the products’ side automatically decomposes according to the following equation,

I3- -> I- + I2

The solution is then titrated with sodium thiosulfate until a pale yellow color is achieved. 1% starch solution indicator is added, making the solution blue, and the solution is further titrated until the colorless endpoint is achieved.

I2 + 2S2O32- -> 2I- + S4O6

2- SAMPLE ANALYSIS Stoichometric Ratio: 1 mmol O2 = 4 mmols S2O3

2-

A water sample filled to the brim in the bottle is opened and manganese sulfate (MnSO4), ammonium hydrogen carbonate (NH4HCO3), and sodium hydroxide with potassium iodide and azide (NaOH w/ KI w/ N3) are added to the solution in that specific order. The following reactions proceed as a result.

Mn2+ + 2OH- -> Mn(OH)2 Mn(OH)2 + O2 + H2O -> Mn(OH)3 or 4MnO(OH)

The reason why there are two possible products for the second reaction is because scientists have yet to confirm which of the two is formed. Upon shaking the solution, a brown precipitate in the form of MnO(OH)2 or Mn(OH)3 is formed, and at this point, the oxygen content is fixed. Concentrated phosphoric acid (H3PO4) is added and the bottle is shaken in order to dissolve the brown precipitate. The following reaction proceeds,

6H+ + 2MnO(OH) + 2I- -> 2Mn2+ + I2 + 4H2O

Once completely dissolved, the solution is titrated with thiosulfate following the same procedure as the standardization process. EXPERIMENTAL CONCEPTS Addition of Manganese Sulfate (MnSO4) Manganese sulfate (MnSO4) is added in order to provide a steady source of Mn2+ ions for the reaction to occur. Following this logic, it is also possible to make use of other sources of the ion, such as manganese chloride (MnCl2). The important thing is to introduce manganese (II) ion to the solution. Be wary though! Look back at the solubility rules for these compounds to make sure they dissolve. Addition of Ammonium Hydrogen Carbonate (NH4HCO3) Ammonium hydrogen carbonate (NH4HCO3) is introduced to prevent the interference of organic compounds in the solution. This is because organic compounds will gradually use up the dissolved oxygen (O2). Addition of Sodium Azide (NaN3) The presence of Sodium azide (NaN3) serves to eliminate the interference of nitrite (NO2-) in the solution. This is because other than thiosulfate, nitrite will also be oxidized by iodine, effectively reducing the volume of titrant used. Addition of Sodium Carbonate (NaCO3)

Addition of this compound to the titrant stabilizes the sodium thiosulfate solution since it decomposes in acidic environment according to the following reaction,

S2O32- + 2H+ -> SO2

-(g) + S(s) + H2O

The carbonate is said to help reform sodium thiosulfate from the constituent products. Addition of Phosphoric Acid (H3PO4) Reason # 1: Acidic Environment for Redox Reaction Explanation: The following equation shows the need for an acidic environment.

6H+ + 2MnO(OH) + 2I- -> 2Mn2+ + I2 + 4H2O

Reason # 2: Dissolution of Precipitate Explanation: A pH value from 1 to 2.5 is needed in order to dissolve the MnO(OH) precipitate into the solution. Reason # 3: Inactivate Fe3+ Ions Explanation: Ferric ions are reduced to ferrous ions, which can be accounted for and consume 0.14g of oxygen in according to the following reaction.

Fe2+ + O2- -> FeO 1% Starch Indicator The need for freshly prepared starch is due to the hydrolysis that starch can undergo according to the following reaction.

Starch -> Amylose + Amylopectin Delay of starch addition, however, is in order to prevent the stabilization of the I2-starch complex. The reason is that starch possesses a helical structure and once the I2 is introduced, it is “encaged” within the helix. Should this happen, achieving the endpoint will be difficult. Iodometry vs. Iodimetry

Redox Method Iodometry Iodimetry

Titrant Sodium Thiosulfate Iodine

Analyte Iodine Sodium Thiosulfate

The only real difference here is the choice of titrant and analyte. CAUSE AND EFFECT Addition of Sulfuric Acid before Potassium Iodide Addition of sulfuric acid before potassium iodide will result into the formation of iodic acid according to the following reaction.

IO3- + H+ -> HIO3

This causes the volume of thiosulfate to decrease, the molarity of thiosulfate to increase and the calculated dissolved oxygen content to increase. This is because HIO3

- is a weak acid, meaning the dissociation into constituent ions is incomplete; thus, only some of the IO3

- will be titrated. Lack of Shaking Cause: After addition of NaOH, bottle was not shaken well. Parameter: Dissolved Oxygen Concentration Effect: Decrease Explanation: Not all the dissolved oxygen is accounted for, possibly because not all of the oxygen was precipitated into MnO(OH) Cause: After addition of H3PO4, the bottle was not shaken well. Parameter: Dissolved Oxygen Concentration Effect: Decrease Explanation: Similar to the previous question, not all of the dissolved oxygen is accounted for because this time around not all the oxygen was dissolved. No Boiled Distilled Water Cause: Boiled distilled water was not used in standardization of sodium thiosulfate. Parameter: Molarity of Sodium Thiosulfate Effect: Decrease Explanation: If the water isn’t boiled, the formation of carbonic acid (H2CO3) is very possible. Should this happen, this would provide a source of hydrogen ions for the following reaction to occur.

S2O32-

(g) + 2H+ -> S(s) + SO2- + H2O

This will cause the true concentration of S2O3

2- to decrease and the volume of titrant required to increase.

Left Alone Cause: Water sample without pre-treatment is left to stand overnight Parameter: Dissolved Oxygen Content Effect: Indeterminate Explanation: The water sample may contain both heterotrophic and photosynthetic organisms that will perform cellular respiration and photosynthesis. Effect on oxygen cannot be ascertained. Cause: Placed in locker Parameter: Dissolved Oxygen Concentration Effect: Decrease Explanation: Dissolved oxygen will be consumed by heterotrophic organisms for respiration. Cause: MnSO4 is added then left for an hour before KI w/ NaOH w/ Azide is added. Parameter: Dissolved Oxygen Concentration Effect: Decrease Explanation: Manganese is light sensitive and thus undergoes reduction. This causes the volume of titrant used to decrease and consequently the calculated dissolved oxygen concentration to decrease too. Additional Oxygen Sources Cause: Partially filled sample bottle Parameter: Dissolved Oxygen Concentration Effect: Increase Explanation: This is due to the additional oxygen that will dissolve coming from the air space left behind. Cause: Pipette is above water level Parameter: Dissolved Oxygen Concentration Effect: Increase Explanation: Additional oxygen will be incorporated into the solution due to the distance between the pipette and the water sample. CALCULATIONS Molarity of Sodium Thiosulfate Titrant

[ =

( )

Remember that whatever you put for the aliquot factor has to make the numerator smaller than the original since you diluted the solutions.

Parts Per Million (ppm) Dissolved Oxygen

ppm O2 =

(

)

Experiment # 10 DETERMINATION OF ELECTRODE POTENTIALS

INTRODUCTION Electrode Cells: Copper, Zinc, Iron, Chlorine, Bromine and Iodine Electrochemistry is a branch of chemistry which deals with the potentials of electrochemical cells. These are made up of an electrode and a solution, and it is here where reduction-oxidation reactions take place. The aim of the experiment is the measure the electric potentials of individual galvanic and electrolytic cells. The salt bridge is an important component of the set-up in order to maintain electro-neutrality. One of the most important things to note is the relationship of the individual Ecell values of the solutions mentioned. The higher the E⁰cell value, the greater the ability to undergo reduction. The reduction trend is as follows,

Cl > Br > I > Fe > Cu > Zn This means that chlorine more readily undergoes reduction compared to the rest and is thus the most powerful oxidizing agent. Zinc, on the other hand, more readily undergoes oxidation and is the most powerful reducing agent. Galvanic Cells versus Electrolytic Cells

Cell Type GALVANIC CELL ELECTROLYTIC CELL

Spontaneity Spontaneous Nonspontaneous Cathode Positive Negative

Anode Negative Positive

Electron Flow Anode to Cathode Cathode to Anode

No. of Containers Two One Electrodes Used

Half Cell Electrode

Zinc Zinc (Zn)

Copper Copper (Cu)

Iron Graphite (C) Chlorine Graphite (C)

Bromine Graphite (C)

Iodine Graphite (C)

The Electrolysis Part Anode: Chloride, Bromide and Iodide ions Reaction: 2Ha- -> Ha2

+ 2e- (Ha = Halide) Cathode: Water Reaction: 2H2O + 2e- -> 2OH- + H2 (Efferevescence) Here are the visible results in each cell once a current is passed through. Chlorine: Effervescence is evident. Bromine: Yellowish Coloration Iodine: Yellowish Coloration Water: Effervescence CALCULATIONS CASE 1: Standard-State Conditions (1M for solutions, 1 atm for gases, 25⁰ C)

E⁰cell = Ecathode – Eanode CASE 2: Non-Standard Conditions

Ecell = E⁰cell –

log Q

Ecell = E⁰cell –

ln Q

Q here is the ratio of the concentration of products and reactants, each raised to their respective coefficients. CASE 3: Halide Group Enot Cell Calculations This type of problem, most often than not, asks you to solve for either the Ecell or E⁰cell value of the reaction. Thus, the following equation will be employed.

Ecell = E⁰cell –

log

For this type of calculation, you will be needing the following essential parameters: current (A), time(s), molarity of halide solution (M), volume of halide solution (mL).

1) [Ha-] = –

Total Ha- moles = (

) ( )

Unreacted Ha- moles = (

) ( )

2) [Ha2] = (

) ( )

*Number of electrons (n) and moles of electron (x) in the equation will be determined by the reduction half-reaction. Example: Cl2 + 2e- -> 2Cl- has n = 2 and x = 2 for its number and moles of electrons respectively.

Experiment # 11 POTENTIOMETRIC DETERMINATION OF THE PURITY AND DISSOCIATION CONSTANT OF

POTASSIUM HYDROGEN PHTHALATE INTRODUCTION Primary Standard: Potassium Hydrogen Phthalate (KHP) Titrant: Sodium Hydroxide (NaOH) Titrant/Analyte: Potassium Hydrogen Phthalate (KHP) Indicator: None Potentiometric titration is a widespread method employed in the titration of turbid solutions. It is more efficient than indicator-based titrations due to the presence of a pH meter for accurate quantitative measurement of the endpoint. This method makes use of the original, first derivative and second derivative graphs to measure the volume of titrant at equivalence point and pH at half-equivalence point (for pKa measurement). THE EXPERIMENTAL SET-UP The set-up involves the following components: a magnetic stirrer, burette with titrant, and pH meter and probe. This walkthrough will tackle the significance of each one at a time. Magnetic Stirrer The magnetic stirrer is employed in order to prevent disturbance of the system. This is because with the burette and pH probe in place, manual stirring is almost impossible without hitting either of the mentioned equipment. Stirring is particularly important in order to disperse the titrant throughout the solution. Imagine if there were no stirring mechanism. Wouldn’t the titrant concentrate itself on the surface of the solution? If the pH probe is within close proximity, it will read the pH of the titrant instead of the pH of the overall solution, causing a drastic increase in the pH reading. Sodium Hydroxide (NaOH) Titrant Simply put, the potassium hydrogen phthalate (KHP) solution is acidic and thus must be titrated with a basic titrant.

pH Probe Storage The pH probe is an essential and delicate tool in the potentiometric process, so it must be cared for properly. It must be stored in an acidic solution with a pH of 3. Why can’t it be stored in distilled water, you ask? Recall that the glass probe contains a large concentration of ions. The glass probe will be destroyed by the migration of ions from the more concentrated area (glass probe) to the less concentrated area (distilled water). This supply absolutely cannot be diminished else the pH probe will break. Calibration The calibration process of the pH meter involves the use of buffer solutions. These are freshly prepared since they are organic compounds that easily decompose. They can be stored for a maximum of one week before being thrown out. First, the electrode should be rinsed with water then dipped in Buffer 7. Adjust the “Slope Knob” until a reading of 7 is achieved. Rinse the electrode once more, but this time dip it in the Buffer 4 solution. Turn the “Buffer Knob” until a reading of 4 pops up. Limitation and Errors Working pH Range: 0.5 – 12 Acid Error: pH < 0.5 Effect: The pH that flashes on the pH meter will be higher than the true pH. Alkaline Error: pH > 12 Effect: pH that flashes on the pH meter will be lower than the true pH. Parts of the Probe

pH

Re

ad

ing

Volume of NaOH (mL)

The pH probe is made up of several components, two of which are significant to this discussion, namely the indicator and reference electrodes. REFERENCE ELECTRODE: AgCl/Ag This is the wire that is found inside the pH meter probe. A general characteristic of the reference electrode is that its Ecell value is independent of the concentration of ions in the solution. INDICATOR/ION-SELECTIVE ELECTRODE: Glass-Membrane Electrode From the name itself, this electrode is selective towards the H+ ion, with its Ecell value dependent on the ion’s concentration in the solution. LIQUID PORTION OF PH PROBE: KCl This part of the probe bridges the reference and indicator electrodes and acts as a supporting electrolyte. HOLE IN THE GLASS-ELECTRODE MEMBRANE

This connects the reference electrode to the surrounding solution. GRAPHICAL ANALYSIS These graphical trends are for the specific case of the experiment, where a weak acid was titrated with a strong base. Graphs may vary for other combinations. Original Graph

The graph generally appears to be s-shaped, and the equivalence point can be found at the inflection point of the graph.

First Derivative Graph

The graph generally appears to be hill-shaped, and the equivalence point can be found at the absolute maximum point of the graph. Second Derivative Graph The graph appears to be lightning-shaped, and the equivalence point can be found at the x-intercept of the graph. For the graphs of diprotic and other polyprotic acids, the number of inflection points, relative maxima and x-intercepts increase depending on how many times the acid can be deprotonated (loss of H+).

Fir

st D

eri

va

tiv

e

V' of Titrant (mL)

Se

con

d D

eri

va

tiv

e

V'' of Titrant (mL)

The second row of graphs depicts the case of a weak base titrated with a strong acid. Measuring pKa pKa values can be easily determined by taking the Henderson-Hasselbalch equation.

pH = pKa + log

At half-equivalence point, it has been determined that the following relationship is true.

[A-] = [HA] pH = pKa

ADVANTAGES OF POTENTIOMETRY OVER INDICATOR METHOD Benefit # 1: Inexpensive Explanation: The pH meter, although initially expensive, is a good investment since it can be used for years after initial use and only requires occasional maintenance. On the other hand, indicators must be purchased after stock depletion. Benefit # 2: Accurate Explanation: The exact numerical quantity of the endpoint can be determined through the pH probe. Indicators, however, exhibit a color change over a pH range, which isn’t very accurate. Benefit # 3: Analysis of Turbid (Murky) Solutions Explanation: Since a pH probe is being used, turbid solutions can be titrated and their endpoints can still be determined. Indicators are not suitable in this case due to the possible color interference from the turbid solutions. CALCULATIONS %Purity of KHP Sample

%Purity = (

)

Experiment # 12 DETERMINATION OF COPPER (II) CONCENTRATION BY COLORIMETRIC METHOD

INTRODUCTION Instrument: Single-Beam UV-VIS Spectrophotometer Analyte: Copper (II) Solution The colorimetric method deals specifically with the analysis of colors and their relationship with other properties of the solution. In modern days, a spectrophotometer is used to gauge significant quantities such as absorbance (A) and consequently concentration (M or ppm). EXPERIMENTAL CONCEPTS Parts of the Spectrophotometer

Always remember that the spectrophotometer is comprised of the light source, the wavelength selector (monochromator), cuvette holder, detector and signal processor/readout. Transmittance Transmittance (T) is the ratio of light transmitted through the solution (I) and the incident or original light (Io). When light is fully transmitted, transmittance is one and absorbance is zero. Conversely, when light is completely absorbed, transmittance is zero.

T =

Percent transmittance can be taken from transmittance using the following expression.

%T =

x 100%

The reason why transmittance is not used in graphical representation is because its relationship with concentration is exponential.

Beer’s Law Beer-Lambert’s Law is an equation describing the relationship between absorbance (A) and concentration (M or ppm). It shows that there is a linear relationship between the two and that the equation is given by,

A = abc

a represents the absorptivity constant, which changes into ε or molar absorptivity when concentration is in molarity (M). b, on the other hand, is the path length. It is important to note that the units absorptivity (A) depend on the units of path length (b) and concentration (M or ppm) since absorbance (A) has to be unitless. Limitation of Beer’s Law Real Limitation For results to adhere to Beer’s Law, the concentration of solutions used should be dilute, approximately less than 0.01 M. A higher concentration means molecules are in close proximity with each other. The electrostatic interactions in these solutions cause a change in absorptivity of the solution. Chemical Limitation Solutions must not undergo any kind of chemical reaction such as equilibration, dissociation, precipitation, formation, etc. in the experimental process. This is because the color and consequently the absorbance will change. Instrumental Limitation

Factor # 1: Polychromatic Light

Polychromatic light or light with multiple wavelengths will cause a deviation from linearity described by the image above. As E1 increases and E2 decreases, the graph appears to bend forward. Factor # 2: Stray Light There is always a decrease in the absorbance value when stray light is introduced. The numerical determination of this error is given by the following equation.

A = -log

Ps is the quantification of stray light. Factor # 3: Mismatched Cell

An integral part of the experiment is the use of standardized equipment, particularly of cuvettes. The cuvette used in the auto-zero phase should also be employed for the analysis part. Absorbance increases due to the longer path length that needs to be traversed by the light. The plot then deviates from linearity. The Concept Behind Auto-Zero Auto-zero is performed in order to eliminate the effects of quantities that we do not wish to be part of the absorbance (A) reading. In particular, this means negating the effects of the scattering, reflecting and refracting of light by the cuvette and the absorbance of the ammonia solution. This means that neglecting to perform the auto-zero phase results in an increased absorbance reading.

Significance of Lambda Max Ideal Lambda Max: 615 nanometers Reason # 1: Absorbance per Unit Concentration Explanation: The change in absorbance per unit concentration is greatest at this point. This means that the calibration curve becomes much wider and can accommodate a large working range. There is then a higher probability that the absorbance of the sample is within the calibration curve. Reason # 2: Increased Sensitivity Explanation: Even if the concentration of the sample is small, the instrument is still able to detect it efficiently. Reason # 3: Linear Relationship of A and c Explanation: At lambda max, A and c exhibit a linear relationship. At any other point, there is deviation from linearity. Complementary Color It is vital that one remembers the concept of complementary color. The color wheel can be made a reference for this. Basically, the wavelength that is absorbed by the solution is the complement of the wavelength that is transmitted and received by our eyes. For example, if the solution transmits a blue color to our eyes, it means that the orange wavelength is being absorbed.

The Calibration Curve

A number of things should be considered in the construction of the calibration curve. Note first that the blank solution is not included in the graph. Concentration and absorbance values of the standard solutions are plotted and a line (trendline) is taken. The sample solution’s concentration is then determined from the slope of the graph and related using the following algebraic equation,

y = mx + b

Here, y represents the absorbance (A) of the sample solution. X is then determined by isolating the variable. If the absorbance of the sample solution is higher than the absorbance value of the most concentrated standard, the student may dilute the solution. However, if the absorbance of the sample solution is lower than that of the least concentrated solution, a less concentrated standard should be made and its data should be added to the graph. Solution Preparation Ammonia is added to the solution in order to forward the following reaction,

Cu2+ + 4NH3 -> [Cu(NH3)4]2+

The copper-ammonia complex formed has a more intense color than the copper solution, allowing it to be examined effectively by the spectrophotometer. If there is only a limited amount of ammonia, the hydrolysis of ammonia may proceed and cause the formation of copper hydroxide.

NH3 + H2O <-> NH4+ + OH- Cu2+ + OH- <-> Cu(OH)2

Experiment # 13 DETERMINATION OF TOTAL ION CONCENTRATION USING ION-EXCHANGE CHROMATOGRAPHY

INTRODUCTION Mobile Phase: Distilled Water (H2O) Stationary Phase: Cation-Exchange Resin (Dowex 50) Analyte: Copper (II) Solution Ion-exchange chromatography is a method founded on the concept of ion exchange between two principle sites, the mobile phase and the stationary phase. The stationary phase used in the experiment is a resin. An unknown solution is passed through the resin, where the ion replaces H+/OH- ions that are bound to the binding sites. Measurement is done by titrating the solution (or eluate) that is collected in a beaker at the base of the set-up. The exchange process in the experiment is described by the following reaction.

nRSO3-H+ + Mn+ -> (RSO3)nM + nH+

THE EXPERIMENTAL SET-UP The Resin This is the stationary phase of the set-up. There are generally two types of resins, the cation-exchange (CATEX) resin and the anion-exchange (ANEX) resin. As the name states, a CATEX is used for the analysis of cations, while an ANEX is used for the analysis of anions. The following table lists examples of commercially available resins.

Anion-Exchange Resins (ANEX)

Functional Group Basic (OH-) Functional Group

Resin Name Amberlite IR 410 Dowex 2

Functional Group -NH4+-OH N(CH3)-3OH

ANEX Strength Strong Weak

pH Range 0 - 12 0 – 9

Cation-Exchange Resins (CATEX)

Functional Group Acidic (H+) Functional Group

Resin Name Amberlite IR 400 Dowex 50

Functional Group -COO-H+ SO3-H+ CATEX Strength Weak Strong

pH Range 5 - 14 1-14

In the experiment, Dowex 50 was picked as the stationary phase in the analysis of Cu2+ ions. The pH of the environment was not adjusted anymore due to the large functional pH range possessed by the resin. EXPERIMENTAL CONCEPTS Exchange Capacity The exchange capacity of a resin is given by the following equation.

Exchange Capacity =

What this means is that the large the exchange capacity, the more exchange sites the resin has. For CATEX resins, H+ is used, while with ANEX resins, H+ is replaced with OH- in the equation. Exchange Factors There are a number of factors affecting the efficiency of exchange between the solution and the resin. This is usually measured using the time of elution (time it takes for the ion to reach the beaker) and is governed by two factors. Factor # 1: Ionic Charge Explanation: The larger the charge of the ion being analyzed, the greater its affinity for the resin. This means that compared to other ions of lower charge, an ion of higher charge will take longer to reach the beaker due to its attraction for the resin. Example: Aluminum (Al3+), Barium (Ba2+), Potassium (K+) Parameter: Time of Elution

Al3+ > Ba2+ > K+ (Decreasing Time of Elution) K+ < Ba2+ < Al3+ (Increasing Time of Elution)

This just means that K+ will be the first to be eluted (reach the beaker) due to its relatively small charge, while Al3+ will be the last. Factor # 2: Size Explanation: The larger the size of the ion, the greater the affinity of the ion for the resin. Recall that the overall trend for atomic size increases when moving to the left and down of the periodic table. It is important to note that Francium (Fr+) is then the largest element in the table. Example: Potassium (K+), Sodium (Na+), Cesium (Ce+) Parameter: Time of Elution

Ce+ > Na+ > K+ (Decreasing Time of Elution) K+ < Na+ < Ce+ (Increasing Time of Elution)

This just means that K+ will be the first to be eluted (reach the beaker) due to its relatively small size, while Ce+ will be the last. In a combination problem, first priority is given to the charge of the ion. Once charge has been factored in, size then comes into play. Example: Aluminum (Al3+), Calcium (Ca2+), Potassium (K+), Cesium (Ce+), Sulfate (SO3

-) Additional Condition: Assume that this is a CATEX resin. Parameter: Time of Elution

Al3+ > Ca2+ > Ce+ > K+ > SO3- (Decreasing Time of Elution)

SO3- < K+ < Ce+ < Ca2+ < Al3+ (Increasing Time of Elution)

Note that SO3

- has the smallest time of elution since we are using a CATEX and it is an anion, meaning it simply passes through the resin. Moreover, note that Ce+ is larger than K+ and thus takes significant longer to be eluted. Factor # 3: Concentration At the end of the experiment, the copper (II) ions were washed out of the resin using 6.0 M HCl. How is this possible when copper (II) ions with a charge of +2 have a greater affinity than hydrogen ions of charge +1? The answer lies in the fact that 6.0M was used to wash the resin. This concentration is significantly larger than the copper (II) solution, and this allowed the H+ to replace the copper (II) ions bound to the exchange sites. Factor # 4: Cross-Linking of the Resin Polymer Resin is known to be composed of styrene polymers. In order for polymers to be formed from monomers, cross-linking must occur. The problem here is that a high degree of cross-linking translates to less pores and thus less exchange sites. Conversely, a low degree of cross-linking translates to more pores and more exchange sites.

Factor # 5: Power of Hydrogen (pH) Example: Amino Acid Basic pH: Amino acid becomes an anion. Acidic pH: Amino acid becomes a cation. Factor # 6: 15 Drops/Minute Flow Rate This is done in order to ensure that all the copper (II) ions bind to the exchange sites in the resin. If the rate is too fast, an incomplete exchange process will result and less Cu2+ will bind to the sites. If it is too slow, however, the process is deemed inefficient due to the increased amount of time needed to complete the procedure. METHODOLOGY Resin Soaked in Water

This is done to dissolve water-soluble impurities and to ionize the functional groups of the resin. Water-soluble impurities are troublesome because they can block off access to exchange sites. Through immersion in water, the resin was also allowed to swell, causing exposure of exchange sites.

Resin Soaked in Acid This is done in order to replace any bound ions in the resin and replenish the hydrogen ion (H+) supply. Maintenance of Above-Resin Water Level This is in order to prevent the formation of air pockets, which may restrict access of the ions to the binding sites. These air pockets will decrease the number of available binding sites and the amount of ions bound to them. pH Equality of H2O and Eluate Elution Phase Cause: pH equality is not satisfied Parameter: Cu2+ Concentration Effect: Increase Explanation: The calculated Cu2+ concentration will increase because excess H+ coming from the excess acid (poured to replenish H+) will be added to the eluate and subsequently analyzed. Calculations will lead to a higher resultant Cu2+ from a higher H+ concentration. Titration Phase Cause: pH equality is not satisfied Parameter: Cu2+ Concentration Effect: Decrease Explanation: Cu2+ concentration will decrease because if the pH is not equal, that signifies that not all of the Cu2+ was able to bind with the resin and not all of the H+ is present in the eluate.

CALCULATIONS Molarity of Cu2+ in Unknown Sample

[Cu2+] = (

)