American Journal of Physical Chemistry
Volume 4, Issue 3, June 2015, Pages: 21-29

Use of unmodified orange peel for the adsorption of Cd(II), Pb(II) and Hg(II) ions in aqueous solutions

Egwuatu Chinyelu, Umedum Ngozi, Ochiagha Kate, Ogbugo Sixtus

Department of Pure and Industrial Chemistry, Nnamdi Azikiwe University, P. M. B. 5025 Awka, Anambra State, South Eastern Nigeria

Email address:

(E. Chinyelu)

To cite this article:

Egwuatu Chinyelu, Umedum Ngozi, Ochiagha Kate, Ogbugo Sixtus. Use of Unmodified Orange Peel for the Adsorption of Cd(II), Pb(II) and Hg(II) Ions in Aqueous Solutions.American Journal of Physical Chemistry.Vol.4, No. 3, 2015, pp. 21-29. doi: 10.11648/j.ajpc.20150403.11

Abstract: Ground orange peel (GOP) was used as adsorbent for Cd(II), Pb(II) and Hg(II) ions in aqueous solutions. Several experiments with variation of some factors namely: initial concentration of the heavy metal ions, contact time, temperature and pH were carried out. FT-IR spectrum of GOP showed the presence of different functional groups on its surface.  Equilibrium data were analysed by Langmuir, Freundlich, and Temkin isotherms. Hg (II) and Cd (II) adsorptions were better described by Temkin isotherm while that of Pub (II)was best described Freundlich isotherm. Pseudo second order kinetics with higher R2 values described adsorption in all cases. Adsorption of Hg (II) was maximum at pH 4, while those of Cd (II) and Pub (II) ions were maximum at pH 6. Thermodynamics study showed that the adsorption processes of Pb(II) and Hg(II) ions were  endothermic while that of Cd(II) ions was exothermic in nature.

Keywords: Orange Peels Adsorbent, Adsorption Isotherms, Kinetics, Thermodynamics, pH

1. Introduction

Waste water or aqueous solutions from the food, colouring, paper, paints, carpets, rubber, plastics, cosmetics, textile and other commercial based industries are polluted by heavy metals [1-3]. The presence of very low concentrations of these metals in these effluents (less than 1mg/l for Pb, Cd and Hg) is highly visible and undesirable [3,4]. Heavy metals do not degrade into harmless end products in their metabolism and they are accumulated in the food chain: water--plants--animals-humans; thereby posing the greatest threat to the living organism[5]. Effluents from industries can pollute our plants and thus affect humans greatly. Some brewing industries make use of tanks coated with lead which also can contaminate the drinks which we take. More so some of these industries release heavy metals as gases through their chimneys which when inhaled can be dangerous to our system especially the lungs. Therefore it is advisable to site these industries in non-residential areas [6]. Long term exposure to heavy metals can lead to renal dysfunction, obstructive lung disease in humans and has been linked to lung cancer and bone defects [7]. Emissions of lead into the air have caused considerable pollution. Children are particularly susceptible to lead exposure due to high gastrointestinal uptake and the permeable blood–brain barrier [8]. Cadmium emissions have also increased dramatically during the 20th century because cadmium-containing products are rarely re-cycled, but often dumped together with household waste. Chronic exposure to elevated level of cadmium is known to cause renal dysfunction (Fanconi syndrome), bone degradation (itai–itai syndrome), cancer, hypertension, liver damage, and blood damage [9]. The usage of mercury is widespread in industrial processes and in various products (e.g. batteries, lamps and thermometers). It is also used as amalgam for dental fillings and by the pharmaceutical industry. The presence of mercury in fish, waste water, dental amalgams, vaccine preservatives, and in the atmosphere has made this particular toxic metal an increasing focus for health authorities and interest groups [10,11].

There are numerous methods for the removal of metal ions from aqueous solutions and these include: adsorption, membrane separation processes ion-exchange, floatation, electrodialysis, precipitation, reverse osmosis, coagulation etc.[12]. Adsorption is an important and effective technique but the use of some adsorbents like activated carbon has being seen to be expensive for developing countries[13]. This has encouraged research into discovering materials that are both efficient and cheap for scavenging heavy metal ions in industrial waste water[1]

In this work, orange peels were sourced and used as low-cost adsorbent to remove Hg(II), Cd(II) and Pb(II) ions from aqueous solutions. FT-IR analysis, kinetics, equilibrium, thermodynamics, pH studies were employed to evaluate adsorption parameters.

2. Materials and Methods

2.1. Preparation of Adsorbent Materials and Characterization

The orange peels used were collected from Oja market, Mafoluku, Oshodi, Lagos State, Nigeria. They were sun-dried to remove the moisture and to ease grinding and then ground to fine particles with a laboratory blender, sieved to obtain desired particle size of 300nm and stored in a desiccator for later use.

Functional groups present on the orange peels surface were determined by the FT-IR spectroscopy. The KBr pastille method was used for the orange peels sample preparation. About 2mg of dried orange peels and 200mg KBr were mixed and pulverized to obtain a transparent pellet. The reference measurement was performed with pure KBr and the FT-IR spectra were recorded on an M.530 FT-IR detector in the range of 4000-400 cm-1 with a resolution of 2cm-1.

2.2. Reagents and Metal Solutions

All the primary chemicals used in this study were of analytical grade and used without further purification. The stock solution of 1000mg/l of Cd(II), Pb(II) and Hg(II) ions were prepared by dissolving: 2.743g of Cd(NO3)2.4H2O, 1.60g of Pb(NO3)2 and 1.71g of Hg(NO3)2.H2O respectively in 100ml of deionized water and then made up to in 1000ml in  conical flasks. All required concentrations were prepared by serial dilution of the stock solutions. Solutions of 0.5M NaOH and 0.5M H2SO4 were used for pH adjustment.

2.3. Adsorption Studies Experiments

For the adsorption studies, five beakers were filled with 100ml of each metal ion solution of varying concentrations (10-50mg/l).  5g of ground orange peels(GOP) was added to each beaker. The mixture was thoroughly stirred with a stirrer and allowed to stand for 45mins after which 10ml was drawn out and filtered. The filtrate was analyzed using the Atomic Absorption Spectrophotometer AAS (240 FS AA.). In all cases the amount of ions q adsorbed per unit weight of adsorbent at time t and removal efficiency (R) were calculated as:

q =                                    (1)

Where q is the amount of heavy metal ions adsorbed in mg/l, Co is the initial concentration of the metal ion in mg/l, V is the volume of the solution in litres, and Ce is the equilibrium metal ion concentration in mg/l and W is the mass of adsorbent in grams

The kinetics of adsorption was studied at various times of 5, 10,30,60,90 and 120mins. 5g each of GOP was added into six beakers, each containing 100ml of  40mg/l of the metal ions. Concentration and pH constant were kept constant at room temperature. The mixtures were mechanically stirred and left to stand. 10ml of the mixture was drawn out at the end of each interval and filtered. The filtrate was analysed using the Atomic Absorption Spectrophotometer.

Thermodynamics experiments were carried out at temperatures of 50oC, 70oC and 90oC. 5g of GOP each was added into three different beakers containing 40mg/l of the metal ions solution. The mixtures were vigorously shaken, placed in a thermostat water bath already set at 50oC. These were allowed to stand for one hour after which 10ml of the mixture was drawn from each of the beakers separately, filtered and analysed. The above procedure was repeated at temperatures of 70oC and 90oC.

The effect of pH on adsorption of the metal ions was studied over the pH range of 2-8. 100 ml of 40mg/l of Pb(NO3) solution was measured into different beakers. The pH was adjusted from 2-8 using 0.5M H2SO4 and 0.5M NaOH.  5g of bean husk was added into each beaker and stirred. The mixture was allowed to stand for one hour after which 10ml was drawn out and filtered for analysis. This procedure was repeated using Cd(NO3)2.4H2O and Hg(NO3).H2O solutions.

3. Results and Discussions

3.1. FT-IR Analysis

In order to identify some main characteristics functional groups in GOP (fresh and metal-adsorbed), infra-red analyses were carried out and the spectra are shown in fig.1 (a-d). The FT-IR spectrum of GOP indicated hydroxyl, amino and carbonyl groups were present on its surface. The broad vibration around 3080cm-1-3500cm-1 is indicative of the presence of hydroxyl ( –OH) and amino(-NH2) groups of carboxylic acids and amides on the surface of GOP. After adsorption of Cd(II), Pb(II) and Hg(II) ions it was observed that there is slight shift in both absorption wavelengths: Cd(II):3022cm-1-3775cm‑1, Pb(II):3213cm-1-3599cm-1 and Hg(II):3132cm-1-3759cm-1. The peaks at 1582cm-1–11767cm-1 were attributed to stretching vibration of carbonyl group (–C=O). It was clearly observed that this bands were shifted to higher wavelengths after metal adsorption: Cd(II)= 1637cm-1-1834cm‑1, Pb(II)=1664cm-1-1882cm-1 and Hg(II)=1591cm-1-1826cm-1. This confirmed that deprotonated carbonyl groups were involved in adsorption of Cd(II), Pb(II) and Hg(II)) onto GOP. The band 1082 cm-1 can be assigned to the stretching vibration of C–O. This band was clearly shifted after metal adsorption to higher wavelengths of 1088cm-1forPb(II) and lower wavelengths of 1001cm-1 for Hg(II) and 1005cm-1 for Cd(II).

The shifts in the adsorption peaks generally observed indicates the existence of a metal binding process taking place on the surface of the orange peels.

3.2. Adsorption Kinetics

A kinetic study of adsorption is necessary as it provides the information about the adsorption mechanism which is crucial for the practicality of the process. In this work, two kinetic models were applied in order to establish which of them shows the best fit with experimentally obtained data.

The pseudo first order is frequently used to predict metal adsorption kinetics [14]. The rate law for a pseudo-first-order reaction is given as:  

ln (qe(expt)-qt) = lnCe(theo) - k1t                        (2)

Where k1(min.g/mg) is the rate constant of pseudo first order adsorption, qt is the amount of metal in (mg/g) adsorbed at any time, qe(expt) is the amount of metal in (mg/g) adsorbed at equilibrium time obtained from experiments, qe(theo) is the amount of metal in (mg/g) adsorbed at equilibrium time obtained from theoretical model. Thus the rate constant k1 and qe(theo) can be obtained from the slope and the intercept of the plot of ln (qe-qt) against t (fig 2) respectively.

Fig. 1(a). FT-IR spectrum of GOP

Fig. 1(b). FT-IR spectrum of GOP- Pb (II) ion

Fig. 1(c). FT-IR spectrum of GOP- Hg(II) ions

Fig. 1(d). FT-IR spectra of GOP- Cd (II) ions.

Fig. 2. Pseudo first order plot of adsorption.

The pseudo second order kinetic model may be expressed by the equation[14]:


Where k2 (g/mg min) is the equilibrium rate constant for the pseudo second order adsorption and can be obtained from the plot of t/q against t.

Fig. 3. Pseudo second order plot for the adsorption.

Kinetic parameters for the adsorption of the metal ions onto GOP are tabulated (Table 1). The measure of the fit of two models is verified by the correlation coefficient of determination R2.  Comparing the R2 values for the two models shows that the pseudo second order kinetics model gave a better fit in all cases with  R2values of 1. This implies that 100% of all the observed experimental data was replicated by the second order kinetics.  Pseudo second order kinetics model implies that the predominant process here is chemisorption which involving covalent bonding between the adsorbate and the surface of the adsorbent. Chemisorption is usually restricted to just one layer of molecules on the surface, although it may be followed by additional layers of physically adsorbed molecules [15]. Values of qe(theo)  obtained from the second order kinetics model were also much higher than those obtained from the first order kinetics; further indicating that adsorption processes were better described by the former.

Table 1. Kinetics parameters for the Adsorption of Pb(II),Cd(II) and Hg(II) ions

Metal ions   First order Second order
qe(expt) (mg/g) qe(theo) (mg/g) k1(/min) R2 qe(theo) (mg/g) k2min. g/mg R2
Hg(II) 0.7960 0.0115 0.0244 0.4521 39.84 0.124 1
Pb(II) 0.7727 0.0722 0.0323 0.8519 38.91 0.0291 1
Cd(II) 0.7888 0.0186 0.0193 0.755 39.53 0.09277 1

3.3. Adsorption Isotherms

The adsorption isotherm indicates how molecules of adsorbates are partitioned between the adsorbents and the liquid phase at equilibrium. In this study, the equilibrium data obtained for the adsorption of Pb(II), Cd(II) and Hg(II) ions onto GOP were analysed using the Langmuir, Freundlich and Temkin isotherms models[16-18]

The linear form of Langmuir isotherm equation is given as

qe                                 (4)

The reciprocal of the equation above yields


A plot of 1/qe against 1/Ce gives a straight line graph- figs(4-6) with 1/qm as the intercept and 1/qmk1 as slope. Langmuir isotherm is frequently evaluated by a separation factor, RL, which is expressed as follows

RL                              (6)

Where Co is the initial solute concentration. The value of separation indicates the shape of the isotherm and the type of the adsorption. Considering the RL value, adsorption can be irreversible (RL=0), favourable (0<RL<1) hence (RL=1) or unfavourable (RL>1)[19].

Freundlich isotherm is purely empirical and it best describes the adsorption on heterogeneous surface[20]. The Freundlich isotherm equation is shown below in its linear form:

lnq =lnkF + lnCe                       (7)

kF in mg/l and n are Freundlich constants related to sorption capacity of the adsorbent and energy of adsorption respectively. These constants are evaluated from the plot of lnq versus lnCe figs(7-9).

Temkin isotherm is represented by the Linear equation as follows;


KT and b are Temkin constants, R is universal gas constant (8.314J/K/mol) and T is temperature in Kelvin. A plot of qe vs. lnCe (figs 10-12) gives a slope and an intercept from which b and kT are evaluated[21-22].

The isotherm parameters for the adsorption of Cd (II), Pb(II) and Hg(II) ions onto GOP are given in Table 2. Temkin model described the adsorption of both Hg(II) and Cd(II) best with R2 values of 0.9248 and 0.9128 respectively. The adsorption of Pb(II) was best described by the Freundlich model with R2 value of 0.8822. The RL values were found to be less than one (RL<1) for Pb(II), Cd(II) and Hg(II) on their adsorptions to orange peels shows a favourable adsorption. Maximum adsorption capacity follows the trend Cd(II) > Hg(II)> Pb(II). n values for adsorption were all less than 1, indicating.  KF sorption capacity values according to Freundlich were more than the sorption capacity qm obtained from both Langmuir.

Table 2. Isotherm parameters for adsorption of Pb(II), Cd(II) and Hg(II) ions.

Model Parameters Hg(II) Pb(II) Cd(II)
Langmuir qm (mg/g) 7.4619 5.7644 9.0365
KL (l/mg) 1 1 10.02
R2 0.6959 0.835 0.643
RL 0.342 0.561 0.267
Freundlich KF (mg/g) 10.143 7.285 12.013
N 0.5596 0.7675 0.3592
R2 0.8437 0.8822 0.7192
Temkin R2 0.9248 0.063 0.9128

Fig. 4. Langmuir isotherm plot for Cd(II)

Fig. 5. Langmuir isotherm plot for Pb(II)

Fig. 6. Langmuir isotherm plot for Hg(II)

Fig. 7. Freundlich isotherm plot for Cd(II)

Fig. 8. Freundlich isotherm plot for Pb(II).

Fig. 9. Freundlich isotherm plot for Hg(II)

Fig. 10. Temkin isotherm plot for Cd(II)

Fig. 11. Temkin isotherm plot for Pb(II)

Fig. 12. Temkin isotherm plot for Hg(II)

3.4. The Effect of pH

One of the most important factors that affect the adsorption of metal ions is the pH of the solution, thus the pH affects both the adsorbent and the adsorbate chemistry in solution. The effect of pH on the adsorption of the heavy metal ions onto orange peels was studied at pH 2-8 for initial metal ion concentration of 40mg/l. It is observed that the percentage removal of the adsorption increased with increasing pH from 2-6 and then decreased. Also observed was that the Hg (II) ions removal increased between pH 2 and 4 but then decreased from pH 6 to 8. At low pH, the sorbent surfaces are highly protonated with H+ ions, thus reducing available adsorption sites on which the equally positively charged Pb(II), Cd(II) and Hg(II) ions would have adsorbed to. Consequently, percentage adsorption is reduced. At moderate pH between 4 and 6, the concentration of the hydroxonium ions is fairly reduced, allowing for a considerable quantity of the metal ions. Reduction in percentage adsorption at basic pH of 8 may be as a result of formation of  metallic hydroxide complexes at high pH.

Fig. 13. A plot of pH against % removal of ions

3.5. Effect of Temperature

In designing adsorption systems, one should have an understanding of changes expected to occur and how fast they will take place. The fastness of the reaction can be calculated from the knowledge of kinetic studies, but the change in reaction that can be expected during the process requires the brief idea of thermodynamics parameters.

The thermodynamics parameters that must be considered to determine the process are enthalpy of adsorption (H), free energy change (G), and entropy change (S) due to transfer of unit mole of solute from solution onto the solid –liquid interface.

The important thermodynamic function H is very useful whenever there is a differential change in the system. The negative value of H indicates exothermicity and the positive value of H indicates endothermicity. The change in entropy, ΔS, indicates the randomness of the adsorption process. The parameter G is used to determine the spontaneity and the feasibility of the adsorption process. The value of G is calculated using the equation below [23].

G =   –                                (9)

Gibb’s free energy is also given as

G = -RTlnkd                                  (10)

Where, kd (distribution coefficient or equilibrium constant) is given as

Kd =qe/ce                                      (11)

Where qe is the amount adsorbed per unit weight of the solid and Ce is the equilibrium concentration of solute in solution. Rearranging and making the required substitution yields

ln kd =  –  /RT                            (12)

Table 3. Thermodynamics parameters for the adsorption

Metal ions ΔH  (kJ/mol) ΔS kJ/mol/s ΔG(kJ/mol)
323K 343K 363K
Hg(II) 11.07 -49.22 17.00 17.99 18.97
Pb(II) 13.49 -12.01 17.37 17.61 17.85
Cd(II) -14.92 37.49 -27.03 -27.78 -28.53

Fig. 14. Thermodynamics plot for adsorption

The values of ΔS and ΔH (table 3) were calculated from the intercept and slope from the plot of lnkd versus 1/T(fig.14). The effect of temperature on the adsorption of Cd (II), Hg (II) and Pb(II) ions on orange peels was studied at three different temperatures; 323K,343K and 363K. Adsorption of Hg (II) and Pb(II) ions are endothermic with  change in enthalpies ΔH of 11.07KJ/mol and 13.49KJ/mol respectively while  that of Cd(II) ions is negative -14.92KJ/mol indicating that it is exothermic. It is observed from the table that changes in entropy ΔS for Hg and Pb ions are negative: -49.22KJ/mol and -12.0154KJ/mol respectively. This indicates a decrease in the randomness of the adsorbed species, while that of Cd ions which is positive, 37.49J/mol indicating an increase in the degree of randomness.

The negative values obtained in  ΔG for Cd(II) ions indicate feasible and spontaneous adsorption while the positive values for  Pb and Hg ions obtained from the thermodynamic calculations indicates a less feasible and non- spontaneous adsorption.

4. Conclusion

The present study showed that the surfaces of GOP were effective in the removal of   Pb(II) Hg(II) and Cd(II) ions. The data were better fitted into the second order model. Adsorption with GOP is described by well-known isotherms and is also affected by pH.


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