# Development of a Novel Solid-State Sensor Electrode Based on Titanium Thin Film as an Indicator Electrode in Potentiometric and Conductometric Acid-Base Titration in Aqueous Solution

Authors: Nasser Abu Ghalwa
Journal: Journal of Sensors
 (2012)

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## Abstract

A modified Ti/(SnO2 + Sb2O3) electrode was prepared by thermal deposition on titanium substrate and its use as indicator electrode to potentiometric and conductometric acid-base titration in aqueous solution at 298 K was developed. The E-pH curve is linear with slope of 0.0512 V/dec at 298 K. The standard potential of this electrode, E0, was determined with respect to the SCE as reference electrode. The recovery percentages for potentiometric and conductometric acid-base titration for acetic acid against NaOH were calculated. The cell constant, specific conductance, and the molar conductance with dilution for some common electrolytes were measured.


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## Body

## 1. Introduction

Proton sensitive metal oxides seem to be promising alternatives to the state-of-the-art measuring methods and might overcome some problems of classical hydrogen electrodes and reference electrodes [1].Metal oxide electrodes are becoming alternatives to glass electrodes. They, can be used as acid-base indicator electrodes. However in extreme acidic and basic conditions of the pH scale, great deviations of the Nernstian behavior are observed [2]. Recently, many potentiometric methods developed for measuring the acidity and basicity for aqueous solutions under different conditions were reported [3]. Different metal/metal oxide electrodes have been used as acid-base indicator electrodes. Of those, the most frequently used are the antimony, niobium, copper, tantalum, stainless steel, iridium, and tungsten electrodes [4–6].The glass electrode is by far the most commonly used pH sensor due to its sensitivity, selectivity, stability, and long lifetime [7]. However, it has several disadvantages due to the intrinsic nature of the glass membrane. One is the glass bulb is mechanical not so stable and it breaks very easily so that such electrodes cannot be used for food application. Another drawback is that glass electrodes cannot be designed in planar structures. The drawbacks of glass electrodes have led to intensive research for alternative pH electrodes. As a result, various kinds of potentiometric pH sensors have been developed [8–13].Titanium/titanium-oxide electrode has been prepared and used as indicator electrode in potentiometric acid-base and oxidation-reduction titration in aqueous solution. TheE-pH curve is linear with slope closed to the theoretical value 2.303RT/F (0.059 V at 298 K). This electrode has been tested for potentiometric titration of acetic, propanoic, and phosphoric acids against NaOH in aqueous medium at 298 K [14].Ti/PbO2/SnO2 electrode was prepared by thermal deposition of Sn on titanium substrate and its use as indicator electrode in potentiometric titration in aqueous solution at 298 K was studied. The E-pH curve is linear with slope of 0.0501 V/dec at 298 K. The standard potential of this electrode, E0, is computed as 722.9 mV with respect to saturated calomel electrode as reference electrode [15].Conductometric titrations have a more selective character when acid-base, complexometric or precipitation reactions are explored [16–18], thus widening the range of analytical applications. Pantel and Weisz had developed conductometric and high-frequency conductometric methods for the determination of endpoint of catalytic titrations [19]. A simultaneous conductometric titration method for determination of mixtures of acetic acid, monochloroacetic acid, and trichloroacetic acid based on the multivariate calibration partial least squares was proposed by Ghorbani et al. [20]. Fifteen metal acetates in aqueous medium have been conductometrically titrated with trichloroacetic acid by Gaslini and Nahum [21]. Wu et al. fabricated miniaturized polysilicon electrodes for conductometric sensors and compared them to Pt black and Pt electrode. This conductometric setup was used to detect urea concentration [22]. A new design of a dual, UV photometric-contactless conductometric detector was described by Novotny et al. [23].A conductometric biosensor for nitrite detection was developed and reported by Zhang et al. [24]. A conductometric biosensor for phosphate [25] and ammonium detection [26] was developed based on interdigitated conductometric thin-film planar electrodes.A careful review of the literature spotted no mention of an indicator electrodeinpotentiometric and conductometric acid-base titration in aqueous solution and made the intent to explore the development of a novel solid-state sensor electrode based on titanium thin film as an indicator electrodeforpotentiometric and conductometric acid-base titration in aqueous solution.In the present study the modified electrode, Ti/(SnO2 + Sb2O3), is prepared and used for the first time as an indicator electrode against standard calomel electrode for potentiometric titration. A pair of the above-modified electrode is used for conductometric acid-base and oxidation-reduction titrations in aqueous solutions.


## 2. Experimental

### 2.1. Preparation of Electrodes: Ti/(SnO2 + Sb2O3)

#### 2.1.1. Treatment of Titanium Surface

The pretreatment of titanium surface was carried out according to the procedures suggested by Yang et al. [27] as follows: a titanium sheet was polished on 320-grit paper strips, using water as lubricant, followed by 1 μm siliceous past blasting. The sheet was then degreased in 40% NaOH, cleaned in a hot 1 : 1 mixture of HNO3 and H2SO4, and finally washed with water. The treated surface was immersed in a boiling aqueous solution of oxalic acid, 15% until the TiO2 dissolved. Coating by active metal oxide was carried out immediately to minimize formation of TiO2 [28].


#### 2.1.2. Coating of Mixture of SnO2 and Sb2O3 on Ti/Surface

The thermal deposition of SnO2 on titanium sheet was carried out according to the procedure of Lipp and Pletcher [29] as follows: the pretreated titanium sheet was  immersed for a few minutes in a solution of SnCl4·5H2O, 20% wt and SbCl3, 2% wt in 2-propanol. The excess of alcohol was allowed to evaporate by heating in air at 90°C for 10 min. This stage was then repeated. After two applications, the oxide layer was formed thermally by heating at 500°C for 20 min in a muffle furnace, with a low and continuous flow of oxygen. These stages were repeated until the SnO2 coating reached a loading of 0.8 (±0.01) mg cm−2. It was normal to employ five applications. Finally, the coating was annealed at 500°C for 60 minutes [28].


### 2.2. Chemicals

The chemicals used in potentiometric titrations and preparation the electrode were, CH3COOH, ferrous ammonium sulphate, NaOH, KMnO4, HNO3, H2SO4, NaF, 2-propanol, SnCl4·5H2O, SbCl3, and oxalic acid. The chemicals are of analytical pure grade.


### 2.3. Sensor Design of Potentiometric Cell

#### 2.3.1. Developed Electrode

The Ti/(SnO2 + Sb2O3) electrode was (20 cm × 1 cm) and about 2 mm thick. It was coupled with a saturated calomel electrode (SCE) to complete the electrical circuit and measure the potential difference between these two electrodes using a digital multimeter model 1008 (Kyoritsu, Japan). Potentials were measured with an accuracy of ±5mV. The error in the measurement of the potential, due to liquid-junction potentials in these electrolytes, is estimated to be about 0.005 V [30].


#### 2.3.2. Electrode Calibration in Potentiometric

Ti/(SnO2 + Sb2O3) electrode was calibrated by immersing the electrode and a reference electrode in a solution of known pH at constant temperature (25°C) and applying to them a setting potential which is approximately equal to the steady-state potential that would be observed for that electrode with respect to the reference electrode in a solution of that value of pH.


#### 2.3.3. Experimental Setup

Titration was carried out in a Pyrex glass beaker using 50 mL of the reductant and the base or the oxidant placed in a microburette. The solution in a beaker is stirred by means of a magnetic stirrer. The electrodes (indicator and reference) were dipped slowly into the aqueous solution (acid or reductant). After the steady-state potential was attained, the titration of the acid was carried out by addition of 1 mL of the base to the acidic solution, waiting until the steady potential is established and then measured. The potential variation depends on the type of the base, the progress of neutralization process, and on the initial concentration of the acid to be titrated. The results were reproducible and satisfactory with an accuracy of±5mV for potential measurements. The process of addition of the titrant was continued until the equivalence point was reached.


### 2.4. Sensor Design of Conductometric Cell

#### 2.4.1. Developed Conductometric Cell

Conductometric cells consist of two identical pairs of Ti/(SnO2 + Sb2O3) electrodes 1 cm apart. Each electrode was with dimensions (1 cm × 1 cm) and with about 2 mm thickness. They were coupled with each other to complete the electrical circuit and the conductometric sensors were used to measure the resistance (in ohms) or conductance (the reciprocal of resistance) between these two electrodes.


#### 2.4.2. Conductometric Cell Calibration

Calibration of conductivity cell is usually made by using a conductivity simulator (CM-30 V) (TOA Electronic Ltd., 44CT, 121w, Japan). After calibrating the instrument by simulator, connect the sensor and cross check by using standard solutions for single point as per process range.


#### 2.4.3. Experimental Setup

Titration was carried out in a pyrex glass beaker using 100 mL of the acid and the base placed in a microburette. The solution in a beaker is stirred by means of a magnetic stirrer. The cell was dipped slowly into the aqueous solution. After the steady-state potential was attained, the measuring results obtained.


## 2.1. Preparation of Electrodes: Ti/(SnO2 + Sb2O3)

### 2.1.1. Treatment of Titanium Surface

The pretreatment of titanium surface was carried out according to the procedures suggested by Yang et al. [27] as follows: a titanium sheet was polished on 320-grit paper strips, using water as lubricant, followed by 1 μm siliceous past blasting. The sheet was then degreased in 40% NaOH, cleaned in a hot 1 : 1 mixture of HNO3 and H2SO4, and finally washed with water. The treated surface was immersed in a boiling aqueous solution of oxalic acid, 15% until the TiO2 dissolved. Coating by active metal oxide was carried out immediately to minimize formation of TiO2 [28].


### 2.1.2. Coating of Mixture of SnO2 and Sb2O3 on Ti/Surface

The thermal deposition of SnO2 on titanium sheet was carried out according to the procedure of Lipp and Pletcher [29] as follows: the pretreated titanium sheet was  immersed for a few minutes in a solution of SnCl4·5H2O, 20% wt and SbCl3, 2% wt in 2-propanol. The excess of alcohol was allowed to evaporate by heating in air at 90°C for 10 min. This stage was then repeated. After two applications, the oxide layer was formed thermally by heating at 500°C for 20 min in a muffle furnace, with a low and continuous flow of oxygen. These stages were repeated until the SnO2 coating reached a loading of 0.8 (±0.01) mg cm−2. It was normal to employ five applications. Finally, the coating was annealed at 500°C for 60 minutes [28].


## 2.1.1. Treatment of Titanium Surface

The pretreatment of titanium surface was carried out according to the procedures suggested by Yang et al. [27] as follows: a titanium sheet was polished on 320-grit paper strips, using water as lubricant, followed by 1 μm siliceous past blasting. The sheet was then degreased in 40% NaOH, cleaned in a hot 1 : 1 mixture of HNO3 and H2SO4, and finally washed with water. The treated surface was immersed in a boiling aqueous solution of oxalic acid, 15% until the TiO2 dissolved. Coating by active metal oxide was carried out immediately to minimize formation of TiO2 [28].


## 2.1.2. Coating of Mixture of SnO2 and Sb2O3 on Ti/Surface

The thermal deposition of SnO2 on titanium sheet was carried out according to the procedure of Lipp and Pletcher [29] as follows: the pretreated titanium sheet was  immersed for a few minutes in a solution of SnCl4·5H2O, 20% wt and SbCl3, 2% wt in 2-propanol. The excess of alcohol was allowed to evaporate by heating in air at 90°C for 10 min. This stage was then repeated. After two applications, the oxide layer was formed thermally by heating at 500°C for 20 min in a muffle furnace, with a low and continuous flow of oxygen. These stages were repeated until the SnO2 coating reached a loading of 0.8 (±0.01) mg cm−2. It was normal to employ five applications. Finally, the coating was annealed at 500°C for 60 minutes [28].


## 2.2. Chemicals

The chemicals used in potentiometric titrations and preparation the electrode were, CH3COOH, ferrous ammonium sulphate, NaOH, KMnO4, HNO3, H2SO4, NaF, 2-propanol, SnCl4·5H2O, SbCl3, and oxalic acid. The chemicals are of analytical pure grade.


## 2.3. Sensor Design of Potentiometric Cell

### 2.3.1. Developed Electrode

The Ti/(SnO2 + Sb2O3) electrode was (20 cm × 1 cm) and about 2 mm thick. It was coupled with a saturated calomel electrode (SCE) to complete the electrical circuit and measure the potential difference between these two electrodes using a digital multimeter model 1008 (Kyoritsu, Japan). Potentials were measured with an accuracy of ±5mV. The error in the measurement of the potential, due to liquid-junction potentials in these electrolytes, is estimated to be about 0.005 V [30].


### 2.3.2. Electrode Calibration in Potentiometric

Ti/(SnO2 + Sb2O3) electrode was calibrated by immersing the electrode and a reference electrode in a solution of known pH at constant temperature (25°C) and applying to them a setting potential which is approximately equal to the steady-state potential that would be observed for that electrode with respect to the reference electrode in a solution of that value of pH.


### 2.3.3. Experimental Setup

Titration was carried out in a Pyrex glass beaker using 50 mL of the reductant and the base or the oxidant placed in a microburette. The solution in a beaker is stirred by means of a magnetic stirrer. The electrodes (indicator and reference) were dipped slowly into the aqueous solution (acid or reductant). After the steady-state potential was attained, the titration of the acid was carried out by addition of 1 mL of the base to the acidic solution, waiting until the steady potential is established and then measured. The potential variation depends on the type of the base, the progress of neutralization process, and on the initial concentration of the acid to be titrated. The results were reproducible and satisfactory with an accuracy of±5mV for potential measurements. The process of addition of the titrant was continued until the equivalence point was reached.


## 2.3.1. Developed Electrode

The Ti/(SnO2 + Sb2O3) electrode was (20 cm × 1 cm) and about 2 mm thick. It was coupled with a saturated calomel electrode (SCE) to complete the electrical circuit and measure the potential difference between these two electrodes using a digital multimeter model 1008 (Kyoritsu, Japan). Potentials were measured with an accuracy of ±5mV. The error in the measurement of the potential, due to liquid-junction potentials in these electrolytes, is estimated to be about 0.005 V [30].


## 2.3.2. Electrode Calibration in Potentiometric

Ti/(SnO2 + Sb2O3) electrode was calibrated by immersing the electrode and a reference electrode in a solution of known pH at constant temperature (25°C) and applying to them a setting potential which is approximately equal to the steady-state potential that would be observed for that electrode with respect to the reference electrode in a solution of that value of pH.


## 2.3.3. Experimental Setup

Titration was carried out in a Pyrex glass beaker using 50 mL of the reductant and the base or the oxidant placed in a microburette. The solution in a beaker is stirred by means of a magnetic stirrer. The electrodes (indicator and reference) were dipped slowly into the aqueous solution (acid or reductant). After the steady-state potential was attained, the titration of the acid was carried out by addition of 1 mL of the base to the acidic solution, waiting until the steady potential is established and then measured. The potential variation depends on the type of the base, the progress of neutralization process, and on the initial concentration of the acid to be titrated. The results were reproducible and satisfactory with an accuracy of±5mV for potential measurements. The process of addition of the titrant was continued until the equivalence point was reached.


## 2.4. Sensor Design of Conductometric Cell

### 2.4.1. Developed Conductometric Cell

Conductometric cells consist of two identical pairs of Ti/(SnO2 + Sb2O3) electrodes 1 cm apart. Each electrode was with dimensions (1 cm × 1 cm) and with about 2 mm thickness. They were coupled with each other to complete the electrical circuit and the conductometric sensors were used to measure the resistance (in ohms) or conductance (the reciprocal of resistance) between these two electrodes.


### 2.4.2. Conductometric Cell Calibration

Calibration of conductivity cell is usually made by using a conductivity simulator (CM-30 V) (TOA Electronic Ltd., 44CT, 121w, Japan). After calibrating the instrument by simulator, connect the sensor and cross check by using standard solutions for single point as per process range.


### 2.4.3. Experimental Setup

Titration was carried out in a pyrex glass beaker using 100 mL of the acid and the base placed in a microburette. The solution in a beaker is stirred by means of a magnetic stirrer. The cell was dipped slowly into the aqueous solution. After the steady-state potential was attained, the measuring results obtained.


## 2.4.1. Developed Conductometric Cell

Conductometric cells consist of two identical pairs of Ti/(SnO2 + Sb2O3) electrodes 1 cm apart. Each electrode was with dimensions (1 cm × 1 cm) and with about 2 mm thickness. They were coupled with each other to complete the electrical circuit and the conductometric sensors were used to measure the resistance (in ohms) or conductance (the reciprocal of resistance) between these two electrodes.


## 2.4.2. Conductometric Cell Calibration

Calibration of conductivity cell is usually made by using a conductivity simulator (CM-30 V) (TOA Electronic Ltd., 44CT, 121w, Japan). After calibrating the instrument by simulator, connect the sensor and cross check by using standard solutions for single point as per process range.


## 2.4.3. Experimental Setup

Titration was carried out in a pyrex glass beaker using 100 mL of the acid and the base placed in a microburette. The solution in a beaker is stirred by means of a magnetic stirrer. The cell was dipped slowly into the aqueous solution. After the steady-state potential was attained, the measuring results obtained.


## 3. Results and Discussion

### 3.1. The Use of Ti/(SnO2 + Sb2O3) Electrode as an Indicator Electrode in Potentiometric Acid-Base Titration in Aqueous Solution at 298 K

#### 3.1.1. TheE-pH Relation of Ti/(SnO2 + Sb2O3) Electrode

Figure1 shows the change of the open circuit potential, E, of the Ti/(SnO2 + Sb2O3) electrode with pH. The E-pH plot of the Ti/(SnO2 + Sb2O3) electrode fits straight line with slope of 0.0512 V/decade at 298 K.Figure 1
E
-pH relation for Ti/(SnO2 + Sb2O3) electrode at 298 K.This value is close to the magnitude of the term 2.303 RT/F at the corresponding temperature (0.059 V at 298 K) which indicates that the working electrode behaves reversibly and can be used as an indicator electrode for [H+] determination:
(1)Sn+2Sb+5H2O→SnO2+Sb2O3+10H++10e-.
Nernst equation of this reaction can be written as
(2)ETi/SnO2+Sb2O3=E0Ti/(SnO2+Sb2O3)+2.303RT10Flog[SnO2+Sb2O3][H+]10[Sn+2Sb][H2O+]5,
where the activity coefficient for dilute solution equals unit and (Sn + Sb) and (SnO2 + Sb2O3) are solids. H2O in excess, their activities equal unit also
(3)ETi/(SnO2+Sb2O3)=E0Ti/(SnO2+Sb2O3)-2.303RTFpH.This equation is applicable for the reversible behavior of working electrode. From the developed Nernst equation we indicate that working electrodes can be used as pH-indicator. From Figure1 the E0 value of the Ti/(SnO2 + Sb2O3) electrode, that is, the potential at [H+] = 1, is computed as 510 mV relative to the saturated calomel electrode.At high or low pH, the electrode indicates pH less than true value as pH glass electrode, it may be due to damage in electrode or existence of alkali metal ions in solution too.


#### 3.1.2. Potentiometric Acid-Base Titration

Figure2 represents the relation between the volume of 0.1 M NaOH with each potential shift in the titrations of acetic acid. The variation of the Ti/(SnO2 + Sb2O3) electrode potential at 298 K with the different volumes of NaOH followed typical potentiometric titration curves. These curves show slight decrease in potential (to more negative values) with the addition of the titrant.Figure 2
Potentiometric titration of acetic acid with 0.1 M NaOH at 298 K.For locating endpoints, better results are obtained by constructing a plot ofΔE/ΔV against V of titrant. Figure 3 represents ΔE/ΔV against V for the potentiometric titrations of CH3COOH against 0.1 M NaOH. From the plots the values of endpoints are determined. The obtained results are listed in Table 1 for acetic acid.Table 1
The molar amounts A of CH3COOH acid, experimental and theoretical amounts of NaOH, Be, Bt, and recovery percentage (R%) for acid-base potentiometric titrations using Ti/(SnO2 + Sb2O3) indicator electrode.
A(M)
B
e(M) NaOH
B
t(M) NaOH
R%
0.1000
0.100
0.1000
100.00
0.0750
0.082
0.0750
109.30
0.0500
0.045
0.0500
90.24
0.0250
0.027
0.0250
108.00Figure 3
Locating the endpoints for potentiometric titration of acetic acid with 0.1 M NaOH at 298 K.The values of the recovery percentage,R%, for all above titrations are calculated from (4). From the plots the values of endpoints and the values of the recovery percentage, R%, are determined as
(4)R%=BeBt×100%,
where Be is the experimental amount of base and Bt is the theoretical amount of base calculated from the soichiometric equations of neutralization reactions. It is clear from these data that the working electrode can be used as indicator electrode with the satisfactory percentage recovery not less than 96.1% in potentiometric titrations. These differences in the recovery percentage may be attributed to the impurities in the reagents. The values of pKa for different acids can be determined using the method of half neutralization [31]. They are close to the previously reported values listed in Table 2 for the tested acids.Table 2
The molar amounts A of CH3COOH acid and experimental pKa for acid-base titrations using Ti/(SnO2 + Sb2O3) indicator electrode.
A(M)
pKa
0.100
4.75
0.075
4.73
0.050
4.75
0.025
4.76


#### 3.1.3. The Response Time of the Sensor

Figure4 shows the time response curve as the continuous dependence of the potential on time for several step changes of pH. Response time of the Ti/(SnO2 + Sb2O3) pH sensor, in the range of (50–70) seconds, was achieved, which rendered the sensor highly practical.Figure 4
Time response curve of Ti/(SnO2 + Sb2O3) electrode potential at different pH values. Upper curve at pH = 2 and lower for pH = 11.


#### 3.1.4. Effect of Temperature on the Response Characteristics

The Ti/(SnO2 + Sb2O3) pH sensor response was evaluated at different temperature, Figure 5. At lower temperatures like 283 K, the slope of the sensor was about 43.7 mV/decade and the sensor would be used for pH measurements in the range from 2 to 11. However, when the temperature of test solutions was adjusted to 333 K, the slope significantly increased to 56.8 mV/decade and the slope reached the value of 60.4 mV/decade at 353 K.Figure 5
Effect of temperature on slope of Ti/(SnO2 + Sb2O3) pH sensor at 283 K(▲), 333 K(◆) and 353 K(■).Figure6 shows the square of the correlation coefficient (r2) for pH measurements using the solid-state sensor at different temperatures as compared to pH values obtained by a conventional pH electrode (Hanna Instruments HI 1131 pH combination electrode). It was found to increase as the temperature increases, whereas r2 values for measurements at 283 K, 298 K, 333 K, and 353 K were 0.9964, 0.9974, 0.9983, and 0.9994, respectively. This indicates that better results could be obtained at higher temperature. Overall, it can be easily recognized that there is an excellent correlation between the results obtained by the solid-state pH sensor and that of the conventional glass pH electrode.Figure 6
Correlation between the conventional glass pH electrode and Ti/(SnO2 + Sb2O3) pH sensor.


### 3.2. The Use of Ti/(SnO2 + Sb2O3) Electrode as an Indicator Electrode in Conductometric Acid-Base Titration in Aqueous Solution at 298 K

#### 3.2.1. Conductometric Acid-Base Titration

Figure7 shows the relation between the conductance of the pair Ti/(SnO2 + Sb2O3) electrode with 0.1 M NaOH in the titration of different concentrations of acetic acid at 298 K. The conductance of the acetic acid will be low on account of its poor dissociation:
(5)CH3COOH(aq)+[Na+(aq)+OH-(aq)]→CH3COOH(aq)+Na+(aq)+H2O(l).
On plotting the conductance against the volume of the alkali added, the two lines obtained will be as shown in Figure 7. The point of intersection gives the end point.Figure 7
The relation between the conductance of a pair Ti/(SnO2 + Sb2O3) electrode in the titration of different concentration of CH3COOH with 0.1 M NaOH at 298 K.The values of the recovery percentage,R%, for all above titrations are calculated from (4). The calculated values of R% are listed in Table 3. It is clear from these data that the Ti/(SnO2 + Sb2O3) electrode can be used as an indicator electrode with the satisfactory recovery percentage not less than 90.24%. These differences in the recovery percentage may be attributed to the impurities in the reagents. Table 4 represents the relation between the endpoints of CH3COOH for Ti/(SnO2 + Sb2O3) and glass conductive electrode at 298 K.Table 3
The molar amounts A of CH3COOH acid, experimental and theoretical  amounts of NaOH, Be, Bt and recovery percentage (R%) for acid-base conductometric titrations using Ti/(SnO2 + Sb2O3) as indicator electrode.
A(M)
B
e(M) NaOH
B
t(M) NaOH
R%
0.100
0.1000
0.100
100.0
0.075
0.0738
0.075
98.5
0.050
0.0480
0.050
96.1
0.025
0.0255
0.025
102.3Table 4
The endpoints of different concentrations of CH3COOH acid for Ti/(SnO2 + Sb2O3) electrode and glass conductive electrode in aqueous solution at 298 K.
Acid (M)
End point of CH3COOH
G.E*
S.E**
0.0100
21.2
21.3
0.0075
17.5
17.5
0.0050
9.5
9.6
0.0025
5.84
5.86
*Glass electrode,**studied electrode.Table5 shows the values of constant of the conductometric cell obtained on the basis of the calibration with a standard solution of KCl and comparison with glass conductometric electrode.Table 5
The specific conductance of the conductometric cell  on calibration with standard KCl solutions at 273 K.
A(M) KCl
Specific conductance (mScm−1) at 273 K
G.E*
S.E**
0.001
0.148
0.142
0.010
1.413
1.392
0.100
12.88
12.75
*Glass electrode, **studied electrode.


#### 3.2.2. Cell Constant

The conductance obtained will, therefore, have to be multiplied by a certain factor in order to get the specific conductance. This factor is called the cell constant which is determined from the dimensions of the cell.The value of cell constant for Ti/(SnO2 + Sb2O3) electrodes in aqueous solutions at 298 K equals 148.4 m−1. Specific conductance of normal, decinormal, and centinormal solutions of potassium chloride at different temperatures is given in Table 6.Table 6
The relation between the specific conductance of KCl solutions for Ti/(SnO2 + Sb2O3) electrode against the glass conductive electrode at 273 K, 291 K,  and 298 K.
A(M)
Specific conductance (Sm−1) at 273 K
Specific conductance (Sm−1) at 291 K
Specific conductance (Sm−1) at 298 K
G.E*
S.E**
G.E*
S.E**
G.E*
S.E**
1.00
6.531
6.528
9.815
9.793
11.169
11.167
0.10
0.715
0.714
1.119
1.119
1.288
1.288
0.01
0.077
0.077
0.122
0.122
0.141
0.141
*Glass electrode, **studied electrode.


#### 3.2.3. Variation of Molar Conductance with Dilution

Molar conductance of an electrolyte increases with the increase in dilution. This was attributed to an increase in the degree of dissociation of the electrolyte. We define the degree of dissociation as the fraction of the total electrolyte in solution which exists in the form of its ion. In other words, on dilution, the same amount of electrolyte is capable of furnishing a large number of ions. It may be pointed out, however, that the increase in the number of ions by dilution is much less than the increases in the volume of the solution. Therefore, the number of ions per unit volume (e.g., per c.c.) actually decreases. Hence, the specific conductance decreases although the molar conductance increases on progressive dilution. It is seen that in each case molar conductance increases with dilution.The variation of molar conductance with dilution for some common electrolytes is shown in Table7 and the variation of molar conductance of infinite dilution in the case of some common electrolytes is shown in Table 8.Table 7
The relation between the molar conductance,Λm, in units of 10−4 Sm2 mol−1, of some common electrolytes and the glass conductive electrode in aqueous solution at 298 K.
A(M)
HCl
NaCl
KCl
H2SO4
CH3COOH
NH4OH
G.E*
S.E**
G.E*
S.E**
G.E*
S.E**
G.E*
S.E**
G.E*
S.E**
G.E*
S.E**
0.1000
391.5
392.04
107.01
107.13
128.84
128.79
249.78
249.72
5.18
5.15
3.6
3.61
0.0100
412.3
411.92
118.51
118.47
141.13
140.09
336.35
336.28
16.09
15.78
16.29
16.287
0.0010
421.4
420.98
123.7
123.67
146.9
147.005
399.5
399.44
48.18
48.11
49.201
49.217
0.0005 M
422.2
423.001
124.42
123.7
147.9
147.97
412.95
412.82
67.54
67.5
67.68
67.678
*Glass electrode, **studied electrode.Table 8
The relation between the molar conductance,Λm0, in units of 10−4 Sm2 mol−1, at infinite dilution for some common electrolytes and that of glass conductive electrode in aqueous solution at 298 K.
Electrolyte
Λ
m
0
(10−4 Sm2 mol−1)
G.E*
S.E**
HCl
426.30
426.40
NaCl
126.23
126.15
KCl
149.90
148.83
CH3COOH
390.51
389.45
*Glass electrode, **studied electrode.


#### 3.2.4. The Response Time of the Sensor

Figure8 shows the time response curve as the continuous dependence of the conductance on time for several step changes of concentrations of acetic acid at 298 K. It is clear that the response time for the studied conductometric cell is nearly 70 second, which rendered the sensor highly practical.Figure 8
Time response curve of Ti/(SnO2 + Sb2O3) electrode potential at different concentrations of acid. Upper curve at 0.1M and lower at 0.1 M.


## 3.1. The Use of Ti/(SnO2 + Sb2O3) Electrode as an Indicator Electrode in Potentiometric Acid-Base Titration in Aqueous Solution at 298 K

### 3.1.1. TheE-pH Relation of Ti/(SnO2 + Sb2O3) Electrode

Figure1 shows the change of the open circuit potential, E, of the Ti/(SnO2 + Sb2O3) electrode with pH. The E-pH plot of the Ti/(SnO2 + Sb2O3) electrode fits straight line with slope of 0.0512 V/decade at 298 K.Figure 1
E
-pH relation for Ti/(SnO2 + Sb2O3) electrode at 298 K.This value is close to the magnitude of the term 2.303 RT/F at the corresponding temperature (0.059 V at 298 K) which indicates that the working electrode behaves reversibly and can be used as an indicator electrode for [H+] determination:
(1)Sn+2Sb+5H2O→SnO2+Sb2O3+10H++10e-.
Nernst equation of this reaction can be written as
(2)ETi/SnO2+Sb2O3=E0Ti/(SnO2+Sb2O3)+2.303RT10Flog[SnO2+Sb2O3][H+]10[Sn+2Sb][H2O+]5,
where the activity coefficient for dilute solution equals unit and (Sn + Sb) and (SnO2 + Sb2O3) are solids. H2O in excess, their activities equal unit also
(3)ETi/(SnO2+Sb2O3)=E0Ti/(SnO2+Sb2O3)-2.303RTFpH.This equation is applicable for the reversible behavior of working electrode. From the developed Nernst equation we indicate that working electrodes can be used as pH-indicator. From Figure1 the E0 value of the Ti/(SnO2 + Sb2O3) electrode, that is, the potential at [H+] = 1, is computed as 510 mV relative to the saturated calomel electrode.At high or low pH, the electrode indicates pH less than true value as pH glass electrode, it may be due to damage in electrode or existence of alkali metal ions in solution too.


### 3.1.2. Potentiometric Acid-Base Titration

Figure2 represents the relation between the volume of 0.1 M NaOH with each potential shift in the titrations of acetic acid. The variation of the Ti/(SnO2 + Sb2O3) electrode potential at 298 K with the different volumes of NaOH followed typical potentiometric titration curves. These curves show slight decrease in potential (to more negative values) with the addition of the titrant.Figure 2
Potentiometric titration of acetic acid with 0.1 M NaOH at 298 K.For locating endpoints, better results are obtained by constructing a plot ofΔE/ΔV against V of titrant. Figure 3 represents ΔE/ΔV against V for the potentiometric titrations of CH3COOH against 0.1 M NaOH. From the plots the values of endpoints are determined. The obtained results are listed in Table 1 for acetic acid.Table 1
The molar amounts A of CH3COOH acid, experimental and theoretical amounts of NaOH, Be, Bt, and recovery percentage (R%) for acid-base potentiometric titrations using Ti/(SnO2 + Sb2O3) indicator electrode.
A(M)
B
e(M) NaOH
B
t(M) NaOH
R%
0.1000
0.100
0.1000
100.00
0.0750
0.082
0.0750
109.30
0.0500
0.045
0.0500
90.24
0.0250
0.027
0.0250
108.00Figure 3
Locating the endpoints for potentiometric titration of acetic acid with 0.1 M NaOH at 298 K.The values of the recovery percentage,R%, for all above titrations are calculated from (4). From the plots the values of endpoints and the values of the recovery percentage, R%, are determined as
(4)R%=BeBt×100%,
where Be is the experimental amount of base and Bt is the theoretical amount of base calculated from the soichiometric equations of neutralization reactions. It is clear from these data that the working electrode can be used as indicator electrode with the satisfactory percentage recovery not less than 96.1% in potentiometric titrations. These differences in the recovery percentage may be attributed to the impurities in the reagents. The values of pKa for different acids can be determined using the method of half neutralization [31]. They are close to the previously reported values listed in Table 2 for the tested acids.Table 2
The molar amounts A of CH3COOH acid and experimental pKa for acid-base titrations using Ti/(SnO2 + Sb2O3) indicator electrode.
A(M)
pKa
0.100
4.75
0.075
4.73
0.050
4.75
0.025
4.76


### 3.1.3. The Response Time of the Sensor

Figure4 shows the time response curve as the continuous dependence of the potential on time for several step changes of pH. Response time of the Ti/(SnO2 + Sb2O3) pH sensor, in the range of (50–70) seconds, was achieved, which rendered the sensor highly practical.Figure 4
Time response curve of Ti/(SnO2 + Sb2O3) electrode potential at different pH values. Upper curve at pH = 2 and lower for pH = 11.


### 3.1.4. Effect of Temperature on the Response Characteristics

The Ti/(SnO2 + Sb2O3) pH sensor response was evaluated at different temperature, Figure 5. At lower temperatures like 283 K, the slope of the sensor was about 43.7 mV/decade and the sensor would be used for pH measurements in the range from 2 to 11. However, when the temperature of test solutions was adjusted to 333 K, the slope significantly increased to 56.8 mV/decade and the slope reached the value of 60.4 mV/decade at 353 K.Figure 5
Effect of temperature on slope of Ti/(SnO2 + Sb2O3) pH sensor at 283 K(▲), 333 K(◆) and 353 K(■).Figure6 shows the square of the correlation coefficient (r2) for pH measurements using the solid-state sensor at different temperatures as compared to pH values obtained by a conventional pH electrode (Hanna Instruments HI 1131 pH combination electrode). It was found to increase as the temperature increases, whereas r2 values for measurements at 283 K, 298 K, 333 K, and 353 K were 0.9964, 0.9974, 0.9983, and 0.9994, respectively. This indicates that better results could be obtained at higher temperature. Overall, it can be easily recognized that there is an excellent correlation between the results obtained by the solid-state pH sensor and that of the conventional glass pH electrode.Figure 6
Correlation between the conventional glass pH electrode and Ti/(SnO2 + Sb2O3) pH sensor.


## 3.1.1. TheE-pH Relation of Ti/(SnO2 + Sb2O3) Electrode

Figure1 shows the change of the open circuit potential, E, of the Ti/(SnO2 + Sb2O3) electrode with pH. The E-pH plot of the Ti/(SnO2 + Sb2O3) electrode fits straight line with slope of 0.0512 V/decade at 298 K.Figure 1
E
-pH relation for Ti/(SnO2 + Sb2O3) electrode at 298 K.This value is close to the magnitude of the term 2.303 RT/F at the corresponding temperature (0.059 V at 298 K) which indicates that the working electrode behaves reversibly and can be used as an indicator electrode for [H+] determination:
(1)Sn+2Sb+5H2O→SnO2+Sb2O3+10H++10e-.
Nernst equation of this reaction can be written as
(2)ETi/SnO2+Sb2O3=E0Ti/(SnO2+Sb2O3)+2.303RT10Flog[SnO2+Sb2O3][H+]10[Sn+2Sb][H2O+]5,
where the activity coefficient for dilute solution equals unit and (Sn + Sb) and (SnO2 + Sb2O3) are solids. H2O in excess, their activities equal unit also
(3)ETi/(SnO2+Sb2O3)=E0Ti/(SnO2+Sb2O3)-2.303RTFpH.This equation is applicable for the reversible behavior of working electrode. From the developed Nernst equation we indicate that working electrodes can be used as pH-indicator. From Figure1 the E0 value of the Ti/(SnO2 + Sb2O3) electrode, that is, the potential at [H+] = 1, is computed as 510 mV relative to the saturated calomel electrode.At high or low pH, the electrode indicates pH less than true value as pH glass electrode, it may be due to damage in electrode or existence of alkali metal ions in solution too.


## 3.1.2. Potentiometric Acid-Base Titration

Figure2 represents the relation between the volume of 0.1 M NaOH with each potential shift in the titrations of acetic acid. The variation of the Ti/(SnO2 + Sb2O3) electrode potential at 298 K with the different volumes of NaOH followed typical potentiometric titration curves. These curves show slight decrease in potential (to more negative values) with the addition of the titrant.Figure 2
Potentiometric titration of acetic acid with 0.1 M NaOH at 298 K.For locating endpoints, better results are obtained by constructing a plot ofΔE/ΔV against V of titrant. Figure 3 represents ΔE/ΔV against V for the potentiometric titrations of CH3COOH against 0.1 M NaOH. From the plots the values of endpoints are determined. The obtained results are listed in Table 1 for acetic acid.Table 1
The molar amounts A of CH3COOH acid, experimental and theoretical amounts of NaOH, Be, Bt, and recovery percentage (R%) for acid-base potentiometric titrations using Ti/(SnO2 + Sb2O3) indicator electrode.
A(M)
B
e(M) NaOH
B
t(M) NaOH
R%
0.1000
0.100
0.1000
100.00
0.0750
0.082
0.0750
109.30
0.0500
0.045
0.0500
90.24
0.0250
0.027
0.0250
108.00Figure 3
Locating the endpoints for potentiometric titration of acetic acid with 0.1 M NaOH at 298 K.The values of the recovery percentage,R%, for all above titrations are calculated from (4). From the plots the values of endpoints and the values of the recovery percentage, R%, are determined as
(4)R%=BeBt×100%,
where Be is the experimental amount of base and Bt is the theoretical amount of base calculated from the soichiometric equations of neutralization reactions. It is clear from these data that the working electrode can be used as indicator electrode with the satisfactory percentage recovery not less than 96.1% in potentiometric titrations. These differences in the recovery percentage may be attributed to the impurities in the reagents. The values of pKa for different acids can be determined using the method of half neutralization [31]. They are close to the previously reported values listed in Table 2 for the tested acids.Table 2
The molar amounts A of CH3COOH acid and experimental pKa for acid-base titrations using Ti/(SnO2 + Sb2O3) indicator electrode.
A(M)
pKa
0.100
4.75
0.075
4.73
0.050
4.75
0.025
4.76


## 3.1.3. The Response Time of the Sensor

Figure4 shows the time response curve as the continuous dependence of the potential on time for several step changes of pH. Response time of the Ti/(SnO2 + Sb2O3) pH sensor, in the range of (50–70) seconds, was achieved, which rendered the sensor highly practical.Figure 4
Time response curve of Ti/(SnO2 + Sb2O3) electrode potential at different pH values. Upper curve at pH = 2 and lower for pH = 11.


## 3.1.4. Effect of Temperature on the Response Characteristics

The Ti/(SnO2 + Sb2O3) pH sensor response was evaluated at different temperature, Figure 5. At lower temperatures like 283 K, the slope of the sensor was about 43.7 mV/decade and the sensor would be used for pH measurements in the range from 2 to 11. However, when the temperature of test solutions was adjusted to 333 K, the slope significantly increased to 56.8 mV/decade and the slope reached the value of 60.4 mV/decade at 353 K.Figure 5
Effect of temperature on slope of Ti/(SnO2 + Sb2O3) pH sensor at 283 K(▲), 333 K(◆) and 353 K(■).Figure6 shows the square of the correlation coefficient (r2) for pH measurements using the solid-state sensor at different temperatures as compared to pH values obtained by a conventional pH electrode (Hanna Instruments HI 1131 pH combination electrode). It was found to increase as the temperature increases, whereas r2 values for measurements at 283 K, 298 K, 333 K, and 353 K were 0.9964, 0.9974, 0.9983, and 0.9994, respectively. This indicates that better results could be obtained at higher temperature. Overall, it can be easily recognized that there is an excellent correlation between the results obtained by the solid-state pH sensor and that of the conventional glass pH electrode.Figure 6
Correlation between the conventional glass pH electrode and Ti/(SnO2 + Sb2O3) pH sensor.


## 3.2. The Use of Ti/(SnO2 + Sb2O3) Electrode as an Indicator Electrode in Conductometric Acid-Base Titration in Aqueous Solution at 298 K

### 3.2.1. Conductometric Acid-Base Titration

Figure7 shows the relation between the conductance of the pair Ti/(SnO2 + Sb2O3) electrode with 0.1 M NaOH in the titration of different concentrations of acetic acid at 298 K. The conductance of the acetic acid will be low on account of its poor dissociation:
(5)CH3COOH(aq)+[Na+(aq)+OH-(aq)]→CH3COOH(aq)+Na+(aq)+H2O(l).
On plotting the conductance against the volume of the alkali added, the two lines obtained will be as shown in Figure 7. The point of intersection gives the end point.Figure 7
The relation between the conductance of a pair Ti/(SnO2 + Sb2O3) electrode in the titration of different concentration of CH3COOH with 0.1 M NaOH at 298 K.The values of the recovery percentage,R%, for all above titrations are calculated from (4). The calculated values of R% are listed in Table 3. It is clear from these data that the Ti/(SnO2 + Sb2O3) electrode can be used as an indicator electrode with the satisfactory recovery percentage not less than 90.24%. These differences in the recovery percentage may be attributed to the impurities in the reagents. Table 4 represents the relation between the endpoints of CH3COOH for Ti/(SnO2 + Sb2O3) and glass conductive electrode at 298 K.Table 3
The molar amounts A of CH3COOH acid, experimental and theoretical  amounts of NaOH, Be, Bt and recovery percentage (R%) for acid-base conductometric titrations using Ti/(SnO2 + Sb2O3) as indicator electrode.
A(M)
B
e(M) NaOH
B
t(M) NaOH
R%
0.100
0.1000
0.100
100.0
0.075
0.0738
0.075
98.5
0.050
0.0480
0.050
96.1
0.025
0.0255
0.025
102.3Table 4
The endpoints of different concentrations of CH3COOH acid for Ti/(SnO2 + Sb2O3) electrode and glass conductive electrode in aqueous solution at 298 K.
Acid (M)
End point of CH3COOH
G.E*
S.E**
0.0100
21.2
21.3
0.0075
17.5
17.5
0.0050
9.5
9.6
0.0025
5.84
5.86
*Glass electrode,**studied electrode.Table5 shows the values of constant of the conductometric cell obtained on the basis of the calibration with a standard solution of KCl and comparison with glass conductometric electrode.Table 5
The specific conductance of the conductometric cell  on calibration with standard KCl solutions at 273 K.
A(M) KCl
Specific conductance (mScm−1) at 273 K
G.E*
S.E**
0.001
0.148
0.142
0.010
1.413
1.392
0.100
12.88
12.75
*Glass electrode, **studied electrode.


### 3.2.2. Cell Constant

The conductance obtained will, therefore, have to be multiplied by a certain factor in order to get the specific conductance. This factor is called the cell constant which is determined from the dimensions of the cell.The value of cell constant for Ti/(SnO2 + Sb2O3) electrodes in aqueous solutions at 298 K equals 148.4 m−1. Specific conductance of normal, decinormal, and centinormal solutions of potassium chloride at different temperatures is given in Table 6.Table 6
The relation between the specific conductance of KCl solutions for Ti/(SnO2 + Sb2O3) electrode against the glass conductive electrode at 273 K, 291 K,  and 298 K.
A(M)
Specific conductance (Sm−1) at 273 K
Specific conductance (Sm−1) at 291 K
Specific conductance (Sm−1) at 298 K
G.E*
S.E**
G.E*
S.E**
G.E*
S.E**
1.00
6.531
6.528
9.815
9.793
11.169
11.167
0.10
0.715
0.714
1.119
1.119
1.288
1.288
0.01
0.077
0.077
0.122
0.122
0.141
0.141
*Glass electrode, **studied electrode.


### 3.2.3. Variation of Molar Conductance with Dilution

Molar conductance of an electrolyte increases with the increase in dilution. This was attributed to an increase in the degree of dissociation of the electrolyte. We define the degree of dissociation as the fraction of the total electrolyte in solution which exists in the form of its ion. In other words, on dilution, the same amount of electrolyte is capable of furnishing a large number of ions. It may be pointed out, however, that the increase in the number of ions by dilution is much less than the increases in the volume of the solution. Therefore, the number of ions per unit volume (e.g., per c.c.) actually decreases. Hence, the specific conductance decreases although the molar conductance increases on progressive dilution. It is seen that in each case molar conductance increases with dilution.The variation of molar conductance with dilution for some common electrolytes is shown in Table7 and the variation of molar conductance of infinite dilution in the case of some common electrolytes is shown in Table 8.Table 7
The relation between the molar conductance,Λm, in units of 10−4 Sm2 mol−1, of some common electrolytes and the glass conductive electrode in aqueous solution at 298 K.
A(M)
HCl
NaCl
KCl
H2SO4
CH3COOH
NH4OH
G.E*
S.E**
G.E*
S.E**
G.E*
S.E**
G.E*
S.E**
G.E*
S.E**
G.E*
S.E**
0.1000
391.5
392.04
107.01
107.13
128.84
128.79
249.78
249.72
5.18
5.15
3.6
3.61
0.0100
412.3
411.92
118.51
118.47
141.13
140.09
336.35
336.28
16.09
15.78
16.29
16.287
0.0010
421.4
420.98
123.7
123.67
146.9
147.005
399.5
399.44
48.18
48.11
49.201
49.217
0.0005 M
422.2
423.001
124.42
123.7
147.9
147.97
412.95
412.82
67.54
67.5
67.68
67.678
*Glass electrode, **studied electrode.Table 8
The relation between the molar conductance,Λm0, in units of 10−4 Sm2 mol−1, at infinite dilution for some common electrolytes and that of glass conductive electrode in aqueous solution at 298 K.
Electrolyte
Λ
m
0
(10−4 Sm2 mol−1)
G.E*
S.E**
HCl
426.30
426.40
NaCl
126.23
126.15
KCl
149.90
148.83
CH3COOH
390.51
389.45
*Glass electrode, **studied electrode.


### 3.2.4. The Response Time of the Sensor

Figure8 shows the time response curve as the continuous dependence of the conductance on time for several step changes of concentrations of acetic acid at 298 K. It is clear that the response time for the studied conductometric cell is nearly 70 second, which rendered the sensor highly practical.Figure 8
Time response curve of Ti/(SnO2 + Sb2O3) electrode potential at different concentrations of acid. Upper curve at 0.1M and lower at 0.1 M.


## 3.2.1. Conductometric Acid-Base Titration

Figure7 shows the relation between the conductance of the pair Ti/(SnO2 + Sb2O3) electrode with 0.1 M NaOH in the titration of different concentrations of acetic acid at 298 K. The conductance of the acetic acid will be low on account of its poor dissociation:
(5)CH3COOH(aq)+[Na+(aq)+OH-(aq)]→CH3COOH(aq)+Na+(aq)+H2O(l).
On plotting the conductance against the volume of the alkali added, the two lines obtained will be as shown in Figure 7. The point of intersection gives the end point.Figure 7
The relation between the conductance of a pair Ti/(SnO2 + Sb2O3) electrode in the titration of different concentration of CH3COOH with 0.1 M NaOH at 298 K.The values of the recovery percentage,R%, for all above titrations are calculated from (4). The calculated values of R% are listed in Table 3. It is clear from these data that the Ti/(SnO2 + Sb2O3) electrode can be used as an indicator electrode with the satisfactory recovery percentage not less than 90.24%. These differences in the recovery percentage may be attributed to the impurities in the reagents. Table 4 represents the relation between the endpoints of CH3COOH for Ti/(SnO2 + Sb2O3) and glass conductive electrode at 298 K.Table 3
The molar amounts A of CH3COOH acid, experimental and theoretical  amounts of NaOH, Be, Bt and recovery percentage (R%) for acid-base conductometric titrations using Ti/(SnO2 + Sb2O3) as indicator electrode.
A(M)
B
e(M) NaOH
B
t(M) NaOH
R%
0.100
0.1000
0.100
100.0
0.075
0.0738
0.075
98.5
0.050
0.0480
0.050
96.1
0.025
0.0255
0.025
102.3Table 4
The endpoints of different concentrations of CH3COOH acid for Ti/(SnO2 + Sb2O3) electrode and glass conductive electrode in aqueous solution at 298 K.
Acid (M)
End point of CH3COOH
G.E*
S.E**
0.0100
21.2
21.3
0.0075
17.5
17.5
0.0050
9.5
9.6
0.0025
5.84
5.86
*Glass electrode,**studied electrode.Table5 shows the values of constant of the conductometric cell obtained on the basis of the calibration with a standard solution of KCl and comparison with glass conductometric electrode.Table 5
The specific conductance of the conductometric cell  on calibration with standard KCl solutions at 273 K.
A(M) KCl
Specific conductance (mScm−1) at 273 K
G.E*
S.E**
0.001
0.148
0.142
0.010
1.413
1.392
0.100
12.88
12.75
*Glass electrode, **studied electrode.


## 3.2.2. Cell Constant

The conductance obtained will, therefore, have to be multiplied by a certain factor in order to get the specific conductance. This factor is called the cell constant which is determined from the dimensions of the cell.The value of cell constant for Ti/(SnO2 + Sb2O3) electrodes in aqueous solutions at 298 K equals 148.4 m−1. Specific conductance of normal, decinormal, and centinormal solutions of potassium chloride at different temperatures is given in Table 6.Table 6
The relation between the specific conductance of KCl solutions for Ti/(SnO2 + Sb2O3) electrode against the glass conductive electrode at 273 K, 291 K,  and 298 K.
A(M)
Specific conductance (Sm−1) at 273 K
Specific conductance (Sm−1) at 291 K
Specific conductance (Sm−1) at 298 K
G.E*
S.E**
G.E*
S.E**
G.E*
S.E**
1.00
6.531
6.528
9.815
9.793
11.169
11.167
0.10
0.715
0.714
1.119
1.119
1.288
1.288
0.01
0.077
0.077
0.122
0.122
0.141
0.141
*Glass electrode, **studied electrode.


## 3.2.3. Variation of Molar Conductance with Dilution

Molar conductance of an electrolyte increases with the increase in dilution. This was attributed to an increase in the degree of dissociation of the electrolyte. We define the degree of dissociation as the fraction of the total electrolyte in solution which exists in the form of its ion. In other words, on dilution, the same amount of electrolyte is capable of furnishing a large number of ions. It may be pointed out, however, that the increase in the number of ions by dilution is much less than the increases in the volume of the solution. Therefore, the number of ions per unit volume (e.g., per c.c.) actually decreases. Hence, the specific conductance decreases although the molar conductance increases on progressive dilution. It is seen that in each case molar conductance increases with dilution.The variation of molar conductance with dilution for some common electrolytes is shown in Table7 and the variation of molar conductance of infinite dilution in the case of some common electrolytes is shown in Table 8.Table 7
The relation between the molar conductance,Λm, in units of 10−4 Sm2 mol−1, of some common electrolytes and the glass conductive electrode in aqueous solution at 298 K.
A(M)
HCl
NaCl
KCl
H2SO4
CH3COOH
NH4OH
G.E*
S.E**
G.E*
S.E**
G.E*
S.E**
G.E*
S.E**
G.E*
S.E**
G.E*
S.E**
0.1000
391.5
392.04
107.01
107.13
128.84
128.79
249.78
249.72
5.18
5.15
3.6
3.61
0.0100
412.3
411.92
118.51
118.47
141.13
140.09
336.35
336.28
16.09
15.78
16.29
16.287
0.0010
421.4
420.98
123.7
123.67
146.9
147.005
399.5
399.44
48.18
48.11
49.201
49.217
0.0005 M
422.2
423.001
124.42
123.7
147.9
147.97
412.95
412.82
67.54
67.5
67.68
67.678
*Glass electrode, **studied electrode.Table 8
The relation between the molar conductance,Λm0, in units of 10−4 Sm2 mol−1, at infinite dilution for some common electrolytes and that of glass conductive electrode in aqueous solution at 298 K.
Electrolyte
Λ
m
0
(10−4 Sm2 mol−1)
G.E*
S.E**
HCl
426.30
426.40
NaCl
126.23
126.15
KCl
149.90
148.83
CH3COOH
390.51
389.45
*Glass electrode, **studied electrode.


## 3.2.4. The Response Time of the Sensor

Figure8 shows the time response curve as the continuous dependence of the conductance on time for several step changes of concentrations of acetic acid at 298 K. It is clear that the response time for the studied conductometric cell is nearly 70 second, which rendered the sensor highly practical.Figure 8
Time response curve of Ti/(SnO2 + Sb2O3) electrode potential at different concentrations of acid. Upper curve at 0.1M and lower at 0.1 M.


## 4. Conclusion

The recovery percentage for potentiometric and conductometric acid-base titration using Ti/(SnO2 + Sb2O3) as indicator electrode was calculated. The E-pH curve is linear with slope of 0.0515 V/decade for the Ti/(SnO2 + Sb2O3) electrode at 298 K. This value is close to the theoretical value 2.303 RT/F (0.059 V at 298 K). The standard potential of the tested electrode, E0, is computed as 510 mV with respect to SCE as reference electrode.Acetic acid was successfully potentiometric and conductometric titrated with NaOH as titrant in aqueous medium at 298 K. The cell constant for the pair electrode was determined from the dimensions of the cell. The relation between the specific conductance of standard KCl solutions for the cell against the glass conductive electrode at different temperatures was studied. The variation of molar conductance with dilution for some common electrolytes was tested using the modified electrode. Also the variation of molar conductance of infinite dilution in case of some common electrolytes is determined by each conductive cell and the conductive glass electrode.



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*Source: 681826-2012-11-14.xml*