Introduction

Disinfection of drinking water has been routinely carried out since the dawn of the twentieth century to eradicate and inactivate the pathogenic organisms and prevent waterborne diseases (Uyak et al. 2007).

Trihalomethanes (THMs) are formed during the chlorination of water, when chlorine has a chemical reaction with naturally occurring organic matter, mainly humic and fulvic acids that were already present in the water supply (Farren 2003; Paim et al. 2007; Pavon et al. 2008).

THMs are organohalogen compounds; they are named as derivatives of the compound methane. THMs include chloroform (CFM), bromodichloromethane (BDCM), dibromochloromethane (DBCM) and bromoform (BFM) (Farren 2003; Hesham and Mahmoud 2013; Pavon et al. 2008). Research on health effects showed a relationship between THMs exposure and bladder cancer. Recently, THMs were suspected to cause not only cancer, but also liver and kidney damage, retarded fetus growth, birth defects and possibly miscarriage (De Castro et al. 2019; Nieuwenhuijsen et al. 2000; Wang et al. 2007; Wright et al. 2004).

Considerable research has been directed toward determining the variables of importance in the formation of THMs. The occurrence of THMs in treated and distributed drinking water varies according to the quality of the water source and the operations carried out in the treatment plant. The main influential factors are the nature and amount of natural organic matter (mainly humic substances), type and concentration of disinfectant, contact time, pH, water temperature and bromide ions which mainly impact the distribution of the compounds among the four THMs species, and may affect the yields of the reactions. In general, higher THMs concentrations are expected at higher levels of the above-mentioned parameters (Bond et al. 2014; Nikolaou et al. 2004; Rodriguez et al. 2004; Sadiq and Rodriguez 2004; Uyak and Toroz 2007; Westerhoff et al. 2004; Zhang et al. 2010).

Most research was focused on studying the factors affecting the formation of trihalomethanes. Only limited research was studying the rate constant for the formation of trihalomethanes. The objective of this study was to determine the rate constant for the formation of trihalomethanes at different temperatures (10, 20, 30 and 40 °C) and, consequently, determined the thermodynamic parameters of activation.

Materials and Methods

Sampling and analytical methods

Samples were collected from four positions of Khandaq El-Sharqi canal (Itay El-Baroud, Dinshal, Damanhour and El-Zawya). The samples were mixed and the composite sample was stored at 4 °C till needed.

Physical parameters were measured according to the Standard Methods for the Examination of Water and Wastewater (APHA, AWWA, WEF 2005). Temperature and pH were measured using portable HACH multi-parameter. Turbidity was measured by a nephelometric method (S.M.2130B) using a turbidity meter (HACH 2100N, HACH Co).

Bromide and ammonia were measured, respectively, using ion chromatography (Dionex DX600) according to USEPA (1997; Dionex Application Note 141). Ultraviolet absorbance at 254 nm was measured according to USEPA method 415.3 (USEPA 2005).

Chlorination and study factor affecting trihalomethanes

A stock solution of sodium hypochlorite was used for chlorination and standardized by “iodometric method I” according to S.M. 4500-Cl B (APHA, AWWA, WEF 2005).

Raw water samples were treated with 8 mg/L of chlorine at pH 7.95, and then, the samples were placed in an incubator at different temperatures (10, 20, 30 and 40 °C). The reaction was allowed to proceed for 1, 2, 3, 4, 5, 6, 7 and 8 h. At the end of each contact time, the samples were placed in 60-ml screw-capped glass vials equipped with a PTFE-faced silicone septum which contains 1 g of phosphate buffer and 6 mg of ammonium chloride to prevent further formation of THMs after the designated reaction time and THMs was determined.

Analysis of trihalomethanes

THMs species were measured according to USEPA method 551.1 (USEPA 1995). For each sample, 10 ml was removed; the pH was checked in this 10 ml aliquot to verify that it is within a pH range of 4.8–5.5. Liquid–liquid extraction was used for each of the sample vials, 3 ml of methyl-t-butyl ether (MTBE) was added, and then, 10 g of sodium chloride (NaCl) was added. The sample vials were shaken vigorously for 2 min and left for 5 min to allow the water and MTBE phases to separate using a disposable Pasteur pipette; about 1 ml of solvent phase (organic upper layer) was transferred into the 2-ml autosampler vial for injection using gas chromatography (Agilent 7890A) equipped with electron capture detector (ECD), autosampler, injector and capillary column DB-1 (length 30 m, internal diameter 0.25 mm, film thickness 1.0 µm) that was used for identification and quantification of THMs. Helium was the carrier gas and the nitrogen was the makeup gas. The flow rate was set at 24.8 cm/sec linear velocity at 150 °C. The oven temperature program was 35 °C held for 9 min, then a 1 °C per minute increase to 40 °C which was maintained for 3 min, and finally a 6 °C per minute increase until a temperature of 150 °C was reached and held for 1 min.

Kinetic data and thermodynamic parameters for the formation of trihalomethanes

The formation of THMs may be expressed as follows:

$$\frac{{{\text{d}}\left[ {\text{THMs}} \right]}}{{{\text{d}}t}} = k \, \left[ {{\text{Cl}}_{2} } \right]^{n} \left[ {\text{TOC}} \right]^{m}$$

where [THMs] is the concentration of THMs; [Cl2] is the concentration of chlorine dose; [TOC] is the concentration of precursors; n, m are orders of reaction; k is a rate constant (Li and Zhao 2006).

For a pseudo-first-order reaction, the kinetic equation could be expressed in the form (El Dib and Ali 1995):

$$\ln C = \ln C_{\text{o}} - kt$$

where k is the rate constant and Co and C are the concentration at time zero and time t, respectively.

The least square procedure program was used for calculating the activation parameters of the activated complex of the formation of THMs at 25 °C.

Results and discussion

Four water samples (Itay El-Baroud, Dinshal, Damanhour and El-Zawya) were collected from Khandaq El-Sharqi canal at July. The samples were mixed and the composite sample was stored at 4 °C till needed. The general characteristics of the composite water sample are given in Table 1.

Table 1 General characteristics of the composite water sample collecting from Khandaq El-Sharqi canal
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Factors affecting trihalomethanes

Effect of contact time

The formation of TTHMs was investigated at various contact times from 1 h up to 8 h. These reactions were performed at constant pH (7.95), UVA254 (0.155 cm−1) and chlorine dose (8 mg/L). Increases in contact time resulted in an increase in TTHMs formation during each temperature (10, 20, 30 and 40 °C) as illustrated in Fig. 1. This trend is in agreement with previous studies (Abd El-Shafy and Grunwald 2000; El Dib and Ali 1995; Basiouny et al. 2008).

Fig. 1
figure 1

Effect of contact time and temperature on the formation of total trihalomethanes

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This result indicated that the chlorine first reacts with the active group quickly, leading to the quick formation of DBPs in the beginning. As the chlorination continued, both free chlorine and reactive groups decreased, and as a result, the reaction and DBPs formation slowed down (Hong et al. 2013).

Effect of temperature

The formation of TTHMs was investigated at various temperatures (10, 20, 30 and 40 °C). These reactions were performed at constant pH (7.95), UVA254 (0.155 cm−1) and chlorine dose (8 mg/L). Increasing the temperature resulted in an increase in TTHMs formation during each contact time from 1 h up to 8 h as illustrated in Fig. 1. This trend is in agreement with previous studies (Basiouny et al. 2008; Chowdhury et al. 2010; El Dib and Ali 1995; Fooladvand et al. 2011; Hong et al. 2013; Roccaro et al. 2008).

Kinetic data for trihalomethanes

Since chlorination proceeds in the presence of excess chlorine, the reaction was assumed to be a pseudo-first-order reaction, where the product concentration THMs was measured. Hence, the rate was calculated according to the modified equation (El Dib and Ali 1995).

$$\ln \left( {C_{\infty } - C_{t} } \right) = \ln C_{\infty } - kt$$

where C∞ is THMs formed at extended contact time and Ct is the THMs value at time t. A plot of ln (C∞ − Ct) versus to (t) gives a straight line. Rate constants were calculated from such plots, and the values are given in Table 2. The data showed that the values of rate constant kobs given for CFM, BDCM, DBCM and TTHMs approached each other whereas CFM attained a higher value for each temperature under the same experimental conditions, and the trend of the rate of formation is in the order CFM > BDCM > DBCM. The relatively high value of kobs for CFM indicated that chlorine liberated could attach more effectively organics than bromide. This trend is in agreement with previous studies (El Dib and Ali 1995).

Table 2 Pseudo-first-order rate constant (kobs, s−1) for the formation of trihalomethanes at different temperatures and activation energy (Ea, k J mol−1)
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Arrhenius equation was found to quite obey for the formation of THMs as shown in Fig. 2. Since the plots of ln kobs versus the reciprocal of absolute temperature were linear,

$$k = A\exp \left( { - E_{\text{a}} /RT} \right)$$
$$\ln k = \ln A - E_{\text{a}} /RT$$
Fig. 2
figure 2

Plot of ln kobs versus the reciprocal of the absolute temperature for the formation of trihalomethanes

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Values of activation energies were calculated using the least squares procedure program and are given with their standard deviation in Table 2.

Determination of the activation parameters of the activated complex of the formation of trihalomethanes

Thermodynamic parameters of activation for the formation of THMs (ΔH#, ΔS# and ΔG#) were calculated by least square procedure program at 25 °C using the following equations (Galasstan et al. 1941):

$$\Delta H^{\# } = E_{\text{a}} - RT$$
$$\Delta S^{\# } /R = \ln A - \ln \left( {K_{\text{B}} T\,e/h} \right)$$
$$\Delta G^{\# } =\Delta H^{\# } - T\Delta S^{\# }$$

where KB is Boltzmann constant (1.38 × 10−23 JK−1), e is the charge of the electron (2.7185) and h is Planck’s constant (6.626 × 10−34 Js)

The data obtained are given with their standard deviations in Table 3. The data showed that there is no significant change observed in ΔG#, and this is presumably due to the compensation between ΔH# and ΔS# to each other. The highly negative value of ΔS# indicates the formation of highly restricted transition state and causes loss of entropy (Galasstan et al. 1941).

Table 3 Thermodynamic parameters (ΔH#, ΔG# (k J mol−1) and ΔS# (J mol−1K−1)) for the formation of trihalomethanes at 25 °C
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Plot of ΔH# versus ΔS# gave a linear relation enclosed all the points of the trihalomethanes formation indicates that all systems follow the same mechanism, and the slop β (isokinetic temperature) was found to be 259 K, and this value is lower than the experimental temperature (313–333 K), indicating that the system is entropic controlled (Fig. 3) (Ismail 2012).

Fig. 3
figure 3

Plot of enthalpy of activation (∆H#, J mol−1) versus entropy of activation (∆S#, J mol−1K−1)

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Conclusion

The following conclusions are obtained from “Results and discussion” of this study:

  1. 1.

    Contact time and temperature were effected on the formation of TTHMs.

  2. 2.

    Rate of formation of THMs increases as temperature increases and the rate of formation of CFM is greater than BDCM and DBCM which indicated that chlorine liberated could attach more effectively organics than bromide.

  3. 3.

    Rate of formation of THMs was assumed to be a pseudo-first-order reaction, and as the temperature increases, the rate of formation of THMs increases.

  4. 4.

    Thermodynamic parameters of activation (ΔH#, ΔS# and ΔG#) for the formation of THMs were calculated.