# Adsorption and Photocatalytic Kinetics of Visible-Light Response N-Doped TiO2 Nanocatalyst for Indoor Acetaldehyde Removal under Dark and Light Conditions

Authors: Yu-Hao Lin; Chih-Huang Weng; Jing-Hua Tzeng; Yao-Tung Lin
Journal: International Journal of Photoenergy
 (2016)

---

## Abstract

Understanding the removal nature of the indoor volatile organic compounds under realistic environment conditions would give clear guidance for the development of air purification devices. The study investigated the removal of indoor acetaldehyde using visible-light-responsive N-doped TiO2 (N-TiO2) photocatalyst under visible-light irradiation (light) and in the absence of light (dark). The adsorption kinetics of acetaldehyde onto N-TiO2 followed a pseudo-second-order model. The magnitude of acetaldehyde adsorption is proportional to temperature, and the results were fitted to the Langmuir isotherm model. Moreover, the effect of initial acetaldehyde concentration and visible-light intensity on the photooxidation of acetaldehyde was well described by the Langmuir-Hinshelwood model. Results show that the mesoporous N-TiO2 catalyst had a high ability to absorb acetaldehyde in the dark condition, and then acetaldehyde was subsequently photooxidized under visible-light irradiation. The adsorption capacity was found to increase with decreasing temperature. The negative value of Δ
G° and the positive value of Δ
S° indicate that the adsorption of acetaldehyde onto N-TiO2 was a spontaneous process. Finally, a reaction scheme for removal process of indoor acetaldehyde by N-TiO2 was proposed.


---

## Body

## 1. Introduction

Indoor volatile organic compounds (VOCs) and polycyclic aromatic hydrocarbons are especially serious for people who spend much of their time in airtight dwellings [1, 2]. Generally, there are two options to remove the VOCs: capture and destruction. Adsorption is one of the most reliable techniques to remove the VOCs in capture category. Adsorption process is classified into two types, that is, physical adsorption and chemical sorption, based on the interaction type between adsorbate and adsorbent [3]. In the destruction category, VOCs can be destroyed by various oxidation processes. The photocatalytic oxidation is one of the effective destruction techniques for the removal of VOCs [4]. As one of the major VOCs, acetaldehyde (CH3CHO) has been frequently used as model compound for the indoor air contamination studies [5]; meanwhile, acetaldehyde can adversely affect the human cardiovascular system. Therefore, using photocatalytic technique to remove acetaldehyde is of considerable interest.The uniform spreading of TiO2 on a support medium is a means of increasing the photocatalytic reaction area for indoor air purification [6]. Ever since the coating of TiO2 on fiberglass fibers could successfully remove indoor VOCs under UV-light illumination [7], self-cleaning paints or home appliances using photocatalysis over TiO2-coated materials are beginning to be commercialized to remove indoor VOCs. Unfortunately, TiO2 cannot function effectively when irradiated with fluorescent light and then cannot be widely applied to indoor use. Thus, expanding the range of applications of TiO2 has been investigated through attempts to increase the sensitivity of TiO2 to visible light [8, 9]. Recently, visible-light-responsive TiO2, such as nitrogen-doped TiO2 (N-TiO2) and N,S-codoped TiO2, have been developed and successfully applied to oxidize a variety of VOCs completely to CO2 under visible-light irradiation [10–14]. Efforts have also been made to explore the reaction kinetics and reaction intermediate of photocatalytic oxidation [15]. Various kinetic and environmental parameters, such as initial concentration, light intensity, and relative humidity, have been shown affecting the process of visible-light-responsive TiO2 [16, 17]. A detailed study of the adsorption phenomena occurring on TiO2 could provide useful insight into the understanding of the photodegradation reaction in the dark and light conditions. Bao et al. [18] and Venditti et al. [19] investigated the adsorption kinetic of a visible-light-active TiO2 and the sequential photocatalyst photodegradation process of contaminants under visible-light irradiation and in the absence of light. However, the light intensity from the fluorescent light source in the indoor room only works approximately 10 hr in day time, and the indoor air often flows slowly.The objective of this study is to investigate the adsorption and photooxidation kinetics of visible-light-responsive N-TiO2 nanosized photocatalyst for indoor acetaldehyde removal under visible-light irradiation (light condition) and in the absence of light (dark condition). The influence of key parameters, such as initial acetaldehyde concentration, light intensity, and temperature, was investigated on the removal of acetaldehyde under simulated airtight environment.


## 2. Materials and Methods

### 2.1. Preparation and Characterization of N-TiO2

The procedure for the synthesis of N-TiO2 composite procedure was modified from literature [14, 20]. The precursors, ammonium hydroxide, titanium tetraisopropoxide, and absolute ethanol, were mixed under a 4°C water bath. After the hydrolysis and condensation reaction, colloid was centrifuged and then calcined at 500°C temperature to obtain the N-TiO2 composite. Detailed synthesis of N-TiO2 composite procedure refers to the literatures [14]. Size and morphological characterizations of N-TiO2 were determined using a scanning electron microscope (SEM, S2700 Hitachi, Japan). The optical absorption responses of N-TiO2 were obtained using UV-vis (Hitachi, U-3900H, Japan). The specific surface areas of N-TiO2 were measured by the Brunauer-Emmett-Teller nitrogen absorption specific surface area (BET-N2 SSA) method using a surface area analyzer (BET Micromeritics ASAP 2020, USA). The crystal phases of the N-TiO2 were analyzed by XRD (PANalytical X’Pert Pro MRD, USA).


### 2.2. Adsorption and Photooxidation Kinetic Experiments

Kinetic experiments of acetaldehyde adsorption onto N-TiO2 were performed to establish the effect of time on the adsorption process and to quantify the adsorption rate. Experiments were conducted with different temperatures (°C), initial acetaldehyde concentration (ppmv), and visible-light intensity (mW/cm2). The sequential experimental procedure including dark and visible-light stages was as follows. Firstly, the catalyst dosage of 0.2 g N-TiO2 was coated on the bottom of the reactor, and the desired concentration of acetaldehyde was injected into the reactor to conduct the absorption reaction without light irradiation (dark reaction). In the dark reaction, temperature was controlled at 5, 15, 25, 35, and 45°C by placing the reaction into the isothermal incubator. After reaching the thermal absorption equilibrium in the dark stage, the photooxidation experiment was sequentially implemented under the various light intensities and initial acetaldehyde concentrations. The visible-light illumination was provided by three-colored fluorescent lamps (March T5-8W/865, Taiwan), and the lamps combined with a cut-off filter (k
>
400 nm) were vertically placed outside the reactor, above the center of the reactor. Illumination intensity was measured with a luminance meter (ILT1700, International Technologies, USA). The samples were withdrawn at designed time intervals and analyzed using a gas chromatograph (GC, PerkinElmer, Clarus 500, USA) equipped with a flame ionization detector (FID) and a thermal conductivity detector (TCD). Based on the pervious study [17], O2 concentration can be neglected in the photooxidation when O2 concentration is greater than 10%. Because the changes of O2 concentration in indoor environment are not obvious, the impact of O2 concentration on the pollutant removal was neglected.


### 2.3. Mathematical Models

#### 2.3.1. Adsorption Kinetics

Adsorption isotherm experiments in the dark condition were performed to determine the maximum adsorption capacity and thermodynamic parameters. The amount of acetaldehyde adsorbed per unit mass of the adsorbent (q
t, in μmol/g) at time t (hr) was computed using the following expression:(1)
q
t
=
C
0
-
C
t
V
m
, where C
0 (μmol/L) and C
t (μmol/L) are the acetaldehyde concentrations before and after adsorption, respectively, for time t (hr); m (g) is the amount of catalyst; and V is the volume (L) of the reactor. The kinetic data were analyzed using a pseudo-second-order (PSO) model, expressed as(2)
d
q
t
d
t
=
1
q
e
+
k
2
q
e
-
q
t
2
,where q
e (μmol/g) denotes the amounts of acetaldehyde adsorbed on N-TiO2 (μmol/g) at equilibrium, respectively, and k
2 is the PSO rate constant (g/μmol/hr). After integration of (2) and application of boundary conditions q
t
=
0 at t
=
0 and q
t
=
q
t at t
=
t, the integrated form of the equation becomes(3)
1
q
e
-
q
t
=
1
q
e
+
k
2
twhich can also be written as(4)
t
q
t
=
1
B
+
1
q
e
tin the linear form, where B
=
k
2
q
e
2 is the initial sorption rate as t
→
0. If the PSO kinetics is applicable, the plot of t
/
q
t versus t will show a linear relationship, which allows computation of q
e,  k
2 (g/μmol/hr), and B without having to know any parameter beforehand.Temperature effects on the adsorption of acetaldehyde were further investigated. The linear form of the Arrhenius equation is expressed as follows:(5)
ln
⁡
k
2
=
ln
⁡
A
-
E
a
R
T
,where A is the preexponential factor, R (J/mol-K) is the universal gas constant, T (K) is the absolute temperature, and E
a (kJ/mol) is the activation energy.


#### 2.3.2. Adsorption Isotherms

The Langmuir adsorption isotherm was used to describe the acetaldehyde adsorption:(6)
q
e
=
K
L
C
e
q
max
⁡
1
+
K
L
C
e
,where q
e (μmol/g) is the amount of acetaldehyde adsorbed at equilibrium, K
L (L/mol) is the Langmuir adsorption constant, C
e (ppmv) is the acetaldehyde concentration at equilibrium, and q
max
⁡ (μmol/g) is the maximum adsorption capacity. Equation (6) can be rearranged to(7)
C
e
q
e
=
1
K
L
q
max
⁡
+
1
q
max
⁡
C
e
.From the plot of(
C
e
/
q
e
) versus C
e, the values of q
max
⁡ and K
L are determined for the adsorption system.


#### 2.3.3. Kinetics of Photooxidation

The kinetics of acetaldehyde photooxidation by N-TiO2 was described by using the Langmuir-Hinshelwood (L-H) approach [6]. Based on the assumption of L-H model, the oxidative product of acetaldehyde (i.e., CO2) did not influence the reaction rate. Equations for the dependence of kinetics on environmental variables, such as acetaldehyde concentration and light intensity, were obtained and shown in Table 1.Table 1
Kinetic models for the acetaldehyde photooxidation by N-TiO2.
Type
Model
Light intensity
Rate constant
Adsorption equilibrium constant
K
C
H
3
C
H
O
K
C
H
3
C
H
O
X
I
L–H monomolecular
—
k
K
a
C
a
1
+
K
a
C
a
K
a
C
a
X
1
+
K
a
C
a
X
II
L–H light intensity-monomolecular
I
α
k
′
K
a
C
a
1
+
K
a
C
a
K
a
C
a
X
1
+
K
a
C
a
X
I: light intensity (mW/cm2); k: rate constant (μmole/hr); k
′: intrinsic rate constant (μmole/hr); α: order constant; K
a: adsorption equilibrium constant (1/ppmv); C
a: acetaldehyde concentration (ppmv).


## 2.1. Preparation and Characterization of N-TiO2

The procedure for the synthesis of N-TiO2 composite procedure was modified from literature [14, 20]. The precursors, ammonium hydroxide, titanium tetraisopropoxide, and absolute ethanol, were mixed under a 4°C water bath. After the hydrolysis and condensation reaction, colloid was centrifuged and then calcined at 500°C temperature to obtain the N-TiO2 composite. Detailed synthesis of N-TiO2 composite procedure refers to the literatures [14]. Size and morphological characterizations of N-TiO2 were determined using a scanning electron microscope (SEM, S2700 Hitachi, Japan). The optical absorption responses of N-TiO2 were obtained using UV-vis (Hitachi, U-3900H, Japan). The specific surface areas of N-TiO2 were measured by the Brunauer-Emmett-Teller nitrogen absorption specific surface area (BET-N2 SSA) method using a surface area analyzer (BET Micromeritics ASAP 2020, USA). The crystal phases of the N-TiO2 were analyzed by XRD (PANalytical X’Pert Pro MRD, USA).


## 2.2. Adsorption and Photooxidation Kinetic Experiments

Kinetic experiments of acetaldehyde adsorption onto N-TiO2 were performed to establish the effect of time on the adsorption process and to quantify the adsorption rate. Experiments were conducted with different temperatures (°C), initial acetaldehyde concentration (ppmv), and visible-light intensity (mW/cm2). The sequential experimental procedure including dark and visible-light stages was as follows. Firstly, the catalyst dosage of 0.2 g N-TiO2 was coated on the bottom of the reactor, and the desired concentration of acetaldehyde was injected into the reactor to conduct the absorption reaction without light irradiation (dark reaction). In the dark reaction, temperature was controlled at 5, 15, 25, 35, and 45°C by placing the reaction into the isothermal incubator. After reaching the thermal absorption equilibrium in the dark stage, the photooxidation experiment was sequentially implemented under the various light intensities and initial acetaldehyde concentrations. The visible-light illumination was provided by three-colored fluorescent lamps (March T5-8W/865, Taiwan), and the lamps combined with a cut-off filter (k
>
400 nm) were vertically placed outside the reactor, above the center of the reactor. Illumination intensity was measured with a luminance meter (ILT1700, International Technologies, USA). The samples were withdrawn at designed time intervals and analyzed using a gas chromatograph (GC, PerkinElmer, Clarus 500, USA) equipped with a flame ionization detector (FID) and a thermal conductivity detector (TCD). Based on the pervious study [17], O2 concentration can be neglected in the photooxidation when O2 concentration is greater than 10%. Because the changes of O2 concentration in indoor environment are not obvious, the impact of O2 concentration on the pollutant removal was neglected.


## 2.3. Mathematical Models

### 2.3.1. Adsorption Kinetics

Adsorption isotherm experiments in the dark condition were performed to determine the maximum adsorption capacity and thermodynamic parameters. The amount of acetaldehyde adsorbed per unit mass of the adsorbent (q
t, in μmol/g) at time t (hr) was computed using the following expression:(1)
q
t
=
C
0
-
C
t
V
m
, where C
0 (μmol/L) and C
t (μmol/L) are the acetaldehyde concentrations before and after adsorption, respectively, for time t (hr); m (g) is the amount of catalyst; and V is the volume (L) of the reactor. The kinetic data were analyzed using a pseudo-second-order (PSO) model, expressed as(2)
d
q
t
d
t
=
1
q
e
+
k
2
q
e
-
q
t
2
,where q
e (μmol/g) denotes the amounts of acetaldehyde adsorbed on N-TiO2 (μmol/g) at equilibrium, respectively, and k
2 is the PSO rate constant (g/μmol/hr). After integration of (2) and application of boundary conditions q
t
=
0 at t
=
0 and q
t
=
q
t at t
=
t, the integrated form of the equation becomes(3)
1
q
e
-
q
t
=
1
q
e
+
k
2
twhich can also be written as(4)
t
q
t
=
1
B
+
1
q
e
tin the linear form, where B
=
k
2
q
e
2 is the initial sorption rate as t
→
0. If the PSO kinetics is applicable, the plot of t
/
q
t versus t will show a linear relationship, which allows computation of q
e,  k
2 (g/μmol/hr), and B without having to know any parameter beforehand.Temperature effects on the adsorption of acetaldehyde were further investigated. The linear form of the Arrhenius equation is expressed as follows:(5)
ln
⁡
k
2
=
ln
⁡
A
-
E
a
R
T
,where A is the preexponential factor, R (J/mol-K) is the universal gas constant, T (K) is the absolute temperature, and E
a (kJ/mol) is the activation energy.


### 2.3.2. Adsorption Isotherms

The Langmuir adsorption isotherm was used to describe the acetaldehyde adsorption:(6)
q
e
=
K
L
C
e
q
max
⁡
1
+
K
L
C
e
,where q
e (μmol/g) is the amount of acetaldehyde adsorbed at equilibrium, K
L (L/mol) is the Langmuir adsorption constant, C
e (ppmv) is the acetaldehyde concentration at equilibrium, and q
max
⁡ (μmol/g) is the maximum adsorption capacity. Equation (6) can be rearranged to(7)
C
e
q
e
=
1
K
L
q
max
⁡
+
1
q
max
⁡
C
e
.From the plot of(
C
e
/
q
e
) versus C
e, the values of q
max
⁡ and K
L are determined for the adsorption system.


### 2.3.3. Kinetics of Photooxidation

The kinetics of acetaldehyde photooxidation by N-TiO2 was described by using the Langmuir-Hinshelwood (L-H) approach [6]. Based on the assumption of L-H model, the oxidative product of acetaldehyde (i.e., CO2) did not influence the reaction rate. Equations for the dependence of kinetics on environmental variables, such as acetaldehyde concentration and light intensity, were obtained and shown in Table 1.Table 1
Kinetic models for the acetaldehyde photooxidation by N-TiO2.
Type
Model
Light intensity
Rate constant
Adsorption equilibrium constant
K
C
H
3
C
H
O
K
C
H
3
C
H
O
X
I
L–H monomolecular
—
k
K
a
C
a
1
+
K
a
C
a
K
a
C
a
X
1
+
K
a
C
a
X
II
L–H light intensity-monomolecular
I
α
k
′
K
a
C
a
1
+
K
a
C
a
K
a
C
a
X
1
+
K
a
C
a
X
I: light intensity (mW/cm2); k: rate constant (μmole/hr); k
′: intrinsic rate constant (μmole/hr); α: order constant; K
a: adsorption equilibrium constant (1/ppmv); C
a: acetaldehyde concentration (ppmv).


## 2.3.1. Adsorption Kinetics

Adsorption isotherm experiments in the dark condition were performed to determine the maximum adsorption capacity and thermodynamic parameters. The amount of acetaldehyde adsorbed per unit mass of the adsorbent (q
t, in μmol/g) at time t (hr) was computed using the following expression:(1)
q
t
=
C
0
-
C
t
V
m
, where C
0 (μmol/L) and C
t (μmol/L) are the acetaldehyde concentrations before and after adsorption, respectively, for time t (hr); m (g) is the amount of catalyst; and V is the volume (L) of the reactor. The kinetic data were analyzed using a pseudo-second-order (PSO) model, expressed as(2)
d
q
t
d
t
=
1
q
e
+
k
2
q
e
-
q
t
2
,where q
e (μmol/g) denotes the amounts of acetaldehyde adsorbed on N-TiO2 (μmol/g) at equilibrium, respectively, and k
2 is the PSO rate constant (g/μmol/hr). After integration of (2) and application of boundary conditions q
t
=
0 at t
=
0 and q
t
=
q
t at t
=
t, the integrated form of the equation becomes(3)
1
q
e
-
q
t
=
1
q
e
+
k
2
twhich can also be written as(4)
t
q
t
=
1
B
+
1
q
e
tin the linear form, where B
=
k
2
q
e
2 is the initial sorption rate as t
→
0. If the PSO kinetics is applicable, the plot of t
/
q
t versus t will show a linear relationship, which allows computation of q
e,  k
2 (g/μmol/hr), and B without having to know any parameter beforehand.Temperature effects on the adsorption of acetaldehyde were further investigated. The linear form of the Arrhenius equation is expressed as follows:(5)
ln
⁡
k
2
=
ln
⁡
A
-
E
a
R
T
,where A is the preexponential factor, R (J/mol-K) is the universal gas constant, T (K) is the absolute temperature, and E
a (kJ/mol) is the activation energy.


## 2.3.2. Adsorption Isotherms

The Langmuir adsorption isotherm was used to describe the acetaldehyde adsorption:(6)
q
e
=
K
L
C
e
q
max
⁡
1
+
K
L
C
e
,where q
e (μmol/g) is the amount of acetaldehyde adsorbed at equilibrium, K
L (L/mol) is the Langmuir adsorption constant, C
e (ppmv) is the acetaldehyde concentration at equilibrium, and q
max
⁡ (μmol/g) is the maximum adsorption capacity. Equation (6) can be rearranged to(7)
C
e
q
e
=
1
K
L
q
max
⁡
+
1
q
max
⁡
C
e
.From the plot of(
C
e
/
q
e
) versus C
e, the values of q
max
⁡ and K
L are determined for the adsorption system.


## 2.3.3. Kinetics of Photooxidation

The kinetics of acetaldehyde photooxidation by N-TiO2 was described by using the Langmuir-Hinshelwood (L-H) approach [6]. Based on the assumption of L-H model, the oxidative product of acetaldehyde (i.e., CO2) did not influence the reaction rate. Equations for the dependence of kinetics on environmental variables, such as acetaldehyde concentration and light intensity, were obtained and shown in Table 1.Table 1
Kinetic models for the acetaldehyde photooxidation by N-TiO2.
Type
Model
Light intensity
Rate constant
Adsorption equilibrium constant
K
C
H
3
C
H
O
K
C
H
3
C
H
O
X
I
L–H monomolecular
—
k
K
a
C
a
1
+
K
a
C
a
K
a
C
a
X
1
+
K
a
C
a
X
II
L–H light intensity-monomolecular
I
α
k
′
K
a
C
a
1
+
K
a
C
a
K
a
C
a
X
1
+
K
a
C
a
X
I: light intensity (mW/cm2); k: rate constant (μmole/hr); k
′: intrinsic rate constant (μmole/hr); α: order constant; K
a: adsorption equilibrium constant (1/ppmv); C
a: acetaldehyde concentration (ppmv).


## 3. Results and Discussion

### 3.1. Characterization of N-TiO2

Figure1(a) shows the crystal phases of as-prepared N-TiO2. The X-ray diffraction peaks marked with letters “A” correspond to the anatase phase. A major peak of 2 theta at 25.5° corresponds to crystal plane (101) of anatase. A higher peak intensity is indicative of stronger anatase crystallinity; thus, the synthesized N-TiO2 would be expected to have high photoactivity. Figure 1(a) also shows the synthesized N-TiO2 sample had typical peaks of polycrystalline anatase structure, but no dopant related peaks, such as TiN, which was possibly due to either the movement of dopant ions into interstitial positions of the TiO2 structure or the ion concentration being too low to be detected [21]. Figures 1(b) and 1(c) show the surface morphology of N-TiO2; the SEM measurement result indicates that single N-TiO2 particle in agglomerates exhibited uniform spherical shape with a size range from 20 to 50 nm.Figure 1
Characteristic of N-TiO2 using (a) XRD, (b) SEM, (c), BET, and (d) UV-vis.
(a)
(b)
(c)
(d)Figure1(c) demonstrated that the N-TiO2 exhibited a type IV isotherm and a type H2 hysteresis loop at lower relative pressure region, which are typical characteristics of mesoporous structure with ink bottle pores. The hysteresis loop at lower relative pressure region (0.4
<
P
/
P
0
<
0.8) and higher relative pressure (0.8
<
P
/
P
0
<
1.0) was attributed to smaller mesopore and larger mesopores, respectively. The BET-N2 SSA, pore size, and pore volume of N-TiO2 catalyst are approximately 45.3 m2/g, 6.9 nm, and 0.13 cm3/g, respectively. Such heterogeneous-type pore is conducive to the adsorption of acetaldehyde. A comparison of optical absorption spectra of the N-TiO2 and commercial TiO2 (Degussa P25) samples was shown in Figure 1(d). The absorption at wavelengths shorter than 400 nm can be assigned to the intrinsic band gap absorption of TiO2. The N-TiO2 samples show a stronger absorption than that of P25 in the visible-light region (>400 nm) and a significant red shift of the absorption edge to a lower energy was observed. This result demonstrated that band gap narrowing has successfully been achieved by the doping of N into the TiO2 lattice. Such visible-light absorption of the N-TiO2 originated from the oxygen vacancy, N interstitial doping, and N substitution in the TiO2 structure [14, 22, 23].Figure2 showed the acetaldehyde removal ability using N-TiO2 photocatalyst under dark and light conditions. The photocatalyst was kept in the dark for at most 4 hr to reach the adsorption equilibrium at different starting acetaldehyde concentrations; this result should not be surprising because N-TiO2 is a mesoporous material. As shown in the inset graph in Figure 2, while the acetaldehyde was gradually photooxidized with the light irradiation time, the concentration of CO2 was simultaneously increased, indicating that the converted acetaldehyde was mineralized to CO2. Hence, the importance of the adsorption and photooxidation kinetics is investigated under the following experimental conditions.Figure 2
Adsorption and photooxidation of acetaldehyde onto the N-TiO2 in the absence of light (dark) and under visible-light irradiation (light). Experimental conditions: [CH3CHO] = 176–1010 ppmv and temperature = 25°C.


### 3.2. Adsorption Kinetics

#### 3.2.1. Effect of Reaction Temperature in Dark Condition

Temperature plays a crucial role in adsorption process. Figure3 shows the extent of acetaldehyde adsorption as a function of contact time and temperature in the dark condition. It was observed that the amount of acetaldehyde adsorbed progressively increased with time and temperature. Equilibrium was achieved in 5 hr for a temperature higher than 5°C. In viewing all these kinetic curves, two phases were observed: (i) the first phase, where a rapid adsorption appeared within 1 hr contact time; (ii) the second phase, where a progressive adsorption occurred thereafter. In the fast adsorption phase, approximately 80% of acetaldehyde was removed. The subsequent slow adsorption stage was mainly caused by external active sites already being occupied by acetaldehyde, as well as by the slow diffusion of acetaldehyde into the pore spaces of N-TiO2. In general, adsorption involves a multistep process where the rate of adsorption is limited by intraparticle diffusion.Figure 3
(a) The fitting of pseudo-second-order kinetic plots for the adsorption of acetaldehyde onto N-TiO2 at various temperatures; the solid lines are the best fittings of PSO to the kinetic data. (b) Arrhenius equation plot for the adsorption of acetaldehyde onto N-TiO2. Experimental conditions: C
H
3
C
H
O
i = 176–1010 ppmv.The kinetic data was analyzed using PSO. The fittings of PSO to the kinetic data were presented as solid line in Figure3. Table 2 summarizes the amount of acetaldehyde adsorbed onto N-TiO2 sample at different temperatures and the corresponding fitting parameters of PSO model. Overall, the kinetic data correlate well with PSO models as judged from the high coefficient of determination values (R
2
>
0.95). The rate constant increased from 0.164 to 0.308 g/μmol/hr as temperature increased from 5 to 35°C. Thus, it was confirmed that an increase in temperature results in an increase in the adsorption rate. An increase in temperature may increase the driving force of diffusion across the external boundary layer and increase the rate of diffusion within the pores. Consequently, more acetaldehyde easily enters the interior pore of N-TiO2. The rate constants, k
2, listed in Table 2 can be related to temperature by Arrhenius equation and were used to determine the activation energy (E
a) of this adsorption process. From the Arrhenius plot of ln
⁡
k
2 versus 1
/
T (graph inset of Figure 3), an E
a of 14 kJ/mol was determined for the adsorption system. Normally a diffusion-controlled process has an E
a value lower than 42 kJ/mol [24]. The obtained E
a value is 14.0 kJ/mol, which indicates the adsorption of acetaldehyde that occurred during the diffusion control process [25].Table 2
Predicted constants of pseudo-second-order kinetic model for the adsorption of acetaldehyde onto N-TiO2 at various temperatures.
Temperature (°C)
Pseudo-second-order model
Activation energy
q
e (μmol/g)
k
2 (g/μmol/hr)
R
2
E
a (kJ/mol)
5
0.860
0.164
0.996
14.0
15
0.687
0.230
0.976
25
0.630
0.245
0.977
35
0.540
0.308
0.957


#### 3.2.2. Adsorption Isotherms

Figure4 shows that the magnitude of acetaldehyde adsorption is proportional to temperature, and the results were fitted to the Langmuir isotherm model (equation (6)). Table 3 lists the Langmuir constants (i.e., q
m
a
x and K
L). The high coefficient of determination values confirms that the equilibrium data could be well represented by the Langmuir isotherm. The changes in the free energy (Δ
G°), enthalpy (Δ
H°), and entropy (Δ
S°) in the adsorption reaction system were calculated from the van’t Hoff plot (graph inset of Figure 4) and are listed in Table 3. The negative value of Δ
H° (−3.42 kJ/mol) indicates that the adsorption mechanism is exothermic in nature and that it was a diffusion-controlled process [26]. Moreover, the small value of Δ
H° implies weak bonding between acetaldehyde and the N-TiO2 surface. The negative value of Δ
G° and the positive value of Δ
S° (3.92 cal/K/mol) indicate that the adsorption of acetaldehyde onto N-TiO2 was a spontaneous process. The positive value of Δ
S° in this study suggests increased randomness at the gas-solid interface with some structural changes in the adsorbate and the adsorbent [27, 28]. In other words, a dissociative adsorption of acetaldehyde occurred possibly on the N-TiO2 surface during adsorption. But the low value of Δ
S° may imply that no remarkable change in entropy occurred during the adsorption of acetaldehyde onto N-TiO2. Normally, adsorption of gases causes a decrease in entropy owing to an orderly arrangement of the gas molecules on a surface. A slight increase in the Δ
G° values from −4.60 to −4.53 kcal/mol with a decrease in temperature from 35 to 5°C indicated that the adsorption process was more favorable and spontaneous at low temperatures.Table 3
Thermodynamic parameters for the adsorption of acetaldehyde onto N-TiO2 at various temperatures.
Temperature (°C)
q
m
a
x (μmol/g)
K
L (L/mol)
R
2
Δ
G
° (kcal/mol)
Δ
H
° (kcal/mol)
Δ
S
° (cal/K/mol)
5
196
0.0017
0.998
−4.53
−3.42
3.92
15
290
0.0018
0.979
−4.59
25
336
0.0022
0.992
−4.57
35
444
0.0036
0.982
−4.60Figure 4
(a) Adsorption isotherms of acetaldehyde adsorption on N-TiO2 at various temperatures. Solid lines are the best fit of Langmuir isotherm. (b) van’t Hoff plot for the adsorption of acetaldehyde onto N-TiO2.


### 3.3. Photooxidation Kinetics

#### 3.3.1. Effect of Acetaldehyde Concentration

The influence of various acetaldehyde concentrations on the rate of acetaldehyde photooxidation is shown in Figure5, and the results correlated well with the L-H rate equation. Increasing the acetaldehyde concentration increases the concentration gradient between the bulk phase and the N-TiO2 surface where the adsorption of the reactant becomes the rate-controlling step. Thus, the higher the acetaldehyde concentration, the faster the reaction rate. According to the L-H model, the rate of the reactant molecular reaction is proportional to the fraction of surface covered by the reactants. Light intensity and temperature were held constant to determine the rate constant (k) (model I in Table 1). If the L-H form is valid, a plot of reciprocal of the degradation rate (1/rate) versus reciprocal of the initial acetaldehyde concentration will be linear. The graph inset of Figure 5 shows the corresponding plot of 1/rate versus 1/[CH3CHO] with good linearity. The red circles in Figure 5 represent a good fit of the experimental data to model I. It is clear that the L–H equation applied to one type of adsorption site which describes the initial acetaldehyde concentration dependencies satisfactorily. The values for k and K
a, generated by model fitting of the experimental data for a reaction temperature of 25°C, were 1.2666 μmol/cm2 and 0.0263 ppmv−1, respectively.Figure 5
The effect of acetaldehyde concentration on the photocatalytic reaction. Experimental conditions: temperature = 25°C,I
=
1.01 mW/cm2.


#### 3.3.2. Effect of Visible-Light Intensity

The N-TiO2 absorbs light with a threshold wavelength that is enough to provide energy to overcome the band gap between the valence band and the conduction band. For N-TiO2, the wavelength of the visible light greater than 400 nm can provide enough energy to overcome the band gap (2.8 eV). At enough activation energy, the electrons will transfer between valence band and conduction band to form electron-hole pairs on the catalyst surface. The electron-hole pairs can trigger the photocatalytic oxidation reactions. Figure 6 depicts the effect of light intensity on the reaction rate of acetaldehyde. The effect of light intensity was examined under the reactant gas stream of 161 ppmv CH3CHO at 25°C. The light intensity used in this experiment was in the range of 0.4–3 mW/cm2. The reaction rates increased with increasing the light intensity, which can generate more photons and electron-hole pairs in the following series of free-radical reactions. The reaction rate constant is a function of light intensity according to a reported relationship [6]. By substituting the above relationship in model I under constant initial concentration of acetaldehyde, the rate expression should be(8)
r
=
k
′
′
I
α
k
′
′
=
k
′
K
a
C
a
1
+
K
a
C
a
.Figure 6
The effect of photon flux on the photocatalytic reaction. Experimental conditions: [CH3CHO] = 161 ppmv and temperature = 25°C.By plottinglog
⁡
(
r
) versus log
⁡
(
I
) (graph inset of Figure 6), a straight line with a slope (α) of 0.111 and intercept (log
⁡
k
′
′) of 0.1112 is obtained and then the rate constant (k
′) is 1.2918 μmol/cm2/hr/(mW/cm2)0.1112. The effect of light intensity on the photooxidation reaction was similar to that reported previously [29]. Photocatalytic reactors can be operated in two regimes with respect to light intensity: a first-order regime where the electron-hole pairs are consumed more rapidly by chemical reactions than by recombination, and a 0.5-order regime where the recombination rate dominates [30]. The present result was consistent with the above phenomenon indicating that the dominant mechanism is the electron-hole pair recombination reaction.To include explicit environmental condition dependence in the rate equation, the following assumptions were made: the Langmuir adsorption constants for monolayer adsorption on a homogeneous surface can be applied on predicting the photooxidation rate of the N-TiO2, and the rate constant (k) follows Arrhenius temperature dependence. With these assumptions, a photooxidation rate of the N-TiO2 is given by model II (Table 1). The constants of model II are summarized in Table 4 and the values calculated are in good agreement with the experimental data, as shown by blue squares in Figure 5. The model II was used to identify the range of two key parameters (i.e., acetaldehyde concentrations and visible-light intensity) that would lead to an enhanced acetaldehyde oxidation rate. As evident from Figure 7, when a photocatalytic device with the N-TiO2 catalyst operated at a maximum acetaldehyde concentration of approximately 250 ppmv, the oxidation rate would approach a constant value. Specifically, oxidation rate of the N-TiO2 photocatalyst increased with elevated acetaldehyde concentration and visible-light intensity.Table 4
Langmuir-Hinshelwood parameters for models I and II.
Parameter
Unit
Type of kinetic model
I
II
k
μmol/h
1.2666
—
K
a
1/ppmv
0.0263
0.0263
α
—
—
0.1110
k
′
μmol/cm2/hr/(mW/cm2)0.1112
—
1.2918Figure 7
Predicted oxidation rate dependence on acetaldehyde concentration and light intensity (model II).


### 3.4. Acetaldehyde Removal Mechanism in Indoor Environment Using the N-TiO2

To illustrate clearly the acetaldehyde removal mechanism of N-TiO2 under the indoor environment, a conceptual reaction scheme of the photocatalytic oxidation of the acetaldehyde using N-TiO2 nanomaterial is proposed (Figure 8). In the dark indoor environment, such as nighttime, the acetaldehyde pollutant will diffuse from air phase to the surface of N-TiO2 catalyst. The acetaldehyde would be absorbed onto N-TiO2 particles, and decreased temperature results in an increase in the adsorption rate. When the visible light was turned on, the N-TiO2 catalyst had shown a stronger absorption in the visible-light region of 400–500 nm wavelength possibly due to the fact that the band gap was narrowed by mixing of nitrogen 2p states with O 2p states on the top of the valence band or a creation of N-induced midgap level [31, 32]; the absorption at wavelength above 500 nm is ascribed to oxygen vacancy. The excited electron (e−) was generated from N-TiO2 and transferred to contact with acetaldehyde. At the same time, the remaining holes also migrate outward to form reactive species and react with acetaldehyde in indoor air. Therefore, the acetaldehyde was photooxidized by N-TiO2 and the photooxidation rate increased with elevated acetaldehyde concentration, but it was less sensitive to the changes of visible-light intensity.Figure 8
Conceptual reaction mechanism of N-TiO2 under an indoor environment for acetaldehyde removal.


## 3.1. Characterization of N-TiO2

Figure1(a) shows the crystal phases of as-prepared N-TiO2. The X-ray diffraction peaks marked with letters “A” correspond to the anatase phase. A major peak of 2 theta at 25.5° corresponds to crystal plane (101) of anatase. A higher peak intensity is indicative of stronger anatase crystallinity; thus, the synthesized N-TiO2 would be expected to have high photoactivity. Figure 1(a) also shows the synthesized N-TiO2 sample had typical peaks of polycrystalline anatase structure, but no dopant related peaks, such as TiN, which was possibly due to either the movement of dopant ions into interstitial positions of the TiO2 structure or the ion concentration being too low to be detected [21]. Figures 1(b) and 1(c) show the surface morphology of N-TiO2; the SEM measurement result indicates that single N-TiO2 particle in agglomerates exhibited uniform spherical shape with a size range from 20 to 50 nm.Figure 1
Characteristic of N-TiO2 using (a) XRD, (b) SEM, (c), BET, and (d) UV-vis.
(a)
(b)
(c)
(d)Figure1(c) demonstrated that the N-TiO2 exhibited a type IV isotherm and a type H2 hysteresis loop at lower relative pressure region, which are typical characteristics of mesoporous structure with ink bottle pores. The hysteresis loop at lower relative pressure region (0.4
<
P
/
P
0
<
0.8) and higher relative pressure (0.8
<
P
/
P
0
<
1.0) was attributed to smaller mesopore and larger mesopores, respectively. The BET-N2 SSA, pore size, and pore volume of N-TiO2 catalyst are approximately 45.3 m2/g, 6.9 nm, and 0.13 cm3/g, respectively. Such heterogeneous-type pore is conducive to the adsorption of acetaldehyde. A comparison of optical absorption spectra of the N-TiO2 and commercial TiO2 (Degussa P25) samples was shown in Figure 1(d). The absorption at wavelengths shorter than 400 nm can be assigned to the intrinsic band gap absorption of TiO2. The N-TiO2 samples show a stronger absorption than that of P25 in the visible-light region (>400 nm) and a significant red shift of the absorption edge to a lower energy was observed. This result demonstrated that band gap narrowing has successfully been achieved by the doping of N into the TiO2 lattice. Such visible-light absorption of the N-TiO2 originated from the oxygen vacancy, N interstitial doping, and N substitution in the TiO2 structure [14, 22, 23].Figure2 showed the acetaldehyde removal ability using N-TiO2 photocatalyst under dark and light conditions. The photocatalyst was kept in the dark for at most 4 hr to reach the adsorption equilibrium at different starting acetaldehyde concentrations; this result should not be surprising because N-TiO2 is a mesoporous material. As shown in the inset graph in Figure 2, while the acetaldehyde was gradually photooxidized with the light irradiation time, the concentration of CO2 was simultaneously increased, indicating that the converted acetaldehyde was mineralized to CO2. Hence, the importance of the adsorption and photooxidation kinetics is investigated under the following experimental conditions.Figure 2
Adsorption and photooxidation of acetaldehyde onto the N-TiO2 in the absence of light (dark) and under visible-light irradiation (light). Experimental conditions: [CH3CHO] = 176–1010 ppmv and temperature = 25°C.


## 3.2. Adsorption Kinetics

### 3.2.1. Effect of Reaction Temperature in Dark Condition

Temperature plays a crucial role in adsorption process. Figure3 shows the extent of acetaldehyde adsorption as a function of contact time and temperature in the dark condition. It was observed that the amount of acetaldehyde adsorbed progressively increased with time and temperature. Equilibrium was achieved in 5 hr for a temperature higher than 5°C. In viewing all these kinetic curves, two phases were observed: (i) the first phase, where a rapid adsorption appeared within 1 hr contact time; (ii) the second phase, where a progressive adsorption occurred thereafter. In the fast adsorption phase, approximately 80% of acetaldehyde was removed. The subsequent slow adsorption stage was mainly caused by external active sites already being occupied by acetaldehyde, as well as by the slow diffusion of acetaldehyde into the pore spaces of N-TiO2. In general, adsorption involves a multistep process where the rate of adsorption is limited by intraparticle diffusion.Figure 3
(a) The fitting of pseudo-second-order kinetic plots for the adsorption of acetaldehyde onto N-TiO2 at various temperatures; the solid lines are the best fittings of PSO to the kinetic data. (b) Arrhenius equation plot for the adsorption of acetaldehyde onto N-TiO2. Experimental conditions: C
H
3
C
H
O
i = 176–1010 ppmv.The kinetic data was analyzed using PSO. The fittings of PSO to the kinetic data were presented as solid line in Figure3. Table 2 summarizes the amount of acetaldehyde adsorbed onto N-TiO2 sample at different temperatures and the corresponding fitting parameters of PSO model. Overall, the kinetic data correlate well with PSO models as judged from the high coefficient of determination values (R
2
>
0.95). The rate constant increased from 0.164 to 0.308 g/μmol/hr as temperature increased from 5 to 35°C. Thus, it was confirmed that an increase in temperature results in an increase in the adsorption rate. An increase in temperature may increase the driving force of diffusion across the external boundary layer and increase the rate of diffusion within the pores. Consequently, more acetaldehyde easily enters the interior pore of N-TiO2. The rate constants, k
2, listed in Table 2 can be related to temperature by Arrhenius equation and were used to determine the activation energy (E
a) of this adsorption process. From the Arrhenius plot of ln
⁡
k
2 versus 1
/
T (graph inset of Figure 3), an E
a of 14 kJ/mol was determined for the adsorption system. Normally a diffusion-controlled process has an E
a value lower than 42 kJ/mol [24]. The obtained E
a value is 14.0 kJ/mol, which indicates the adsorption of acetaldehyde that occurred during the diffusion control process [25].Table 2
Predicted constants of pseudo-second-order kinetic model for the adsorption of acetaldehyde onto N-TiO2 at various temperatures.
Temperature (°C)
Pseudo-second-order model
Activation energy
q
e (μmol/g)
k
2 (g/μmol/hr)
R
2
E
a (kJ/mol)
5
0.860
0.164
0.996
14.0
15
0.687
0.230
0.976
25
0.630
0.245
0.977
35
0.540
0.308
0.957


### 3.2.2. Adsorption Isotherms

Figure4 shows that the magnitude of acetaldehyde adsorption is proportional to temperature, and the results were fitted to the Langmuir isotherm model (equation (6)). Table 3 lists the Langmuir constants (i.e., q
m
a
x and K
L). The high coefficient of determination values confirms that the equilibrium data could be well represented by the Langmuir isotherm. The changes in the free energy (Δ
G°), enthalpy (Δ
H°), and entropy (Δ
S°) in the adsorption reaction system were calculated from the van’t Hoff plot (graph inset of Figure 4) and are listed in Table 3. The negative value of Δ
H° (−3.42 kJ/mol) indicates that the adsorption mechanism is exothermic in nature and that it was a diffusion-controlled process [26]. Moreover, the small value of Δ
H° implies weak bonding between acetaldehyde and the N-TiO2 surface. The negative value of Δ
G° and the positive value of Δ
S° (3.92 cal/K/mol) indicate that the adsorption of acetaldehyde onto N-TiO2 was a spontaneous process. The positive value of Δ
S° in this study suggests increased randomness at the gas-solid interface with some structural changes in the adsorbate and the adsorbent [27, 28]. In other words, a dissociative adsorption of acetaldehyde occurred possibly on the N-TiO2 surface during adsorption. But the low value of Δ
S° may imply that no remarkable change in entropy occurred during the adsorption of acetaldehyde onto N-TiO2. Normally, adsorption of gases causes a decrease in entropy owing to an orderly arrangement of the gas molecules on a surface. A slight increase in the Δ
G° values from −4.60 to −4.53 kcal/mol with a decrease in temperature from 35 to 5°C indicated that the adsorption process was more favorable and spontaneous at low temperatures.Table 3
Thermodynamic parameters for the adsorption of acetaldehyde onto N-TiO2 at various temperatures.
Temperature (°C)
q
m
a
x (μmol/g)
K
L (L/mol)
R
2
Δ
G
° (kcal/mol)
Δ
H
° (kcal/mol)
Δ
S
° (cal/K/mol)
5
196
0.0017
0.998
−4.53
−3.42
3.92
15
290
0.0018
0.979
−4.59
25
336
0.0022
0.992
−4.57
35
444
0.0036
0.982
−4.60Figure 4
(a) Adsorption isotherms of acetaldehyde adsorption on N-TiO2 at various temperatures. Solid lines are the best fit of Langmuir isotherm. (b) van’t Hoff plot for the adsorption of acetaldehyde onto N-TiO2.


## 3.2.1. Effect of Reaction Temperature in Dark Condition

Temperature plays a crucial role in adsorption process. Figure3 shows the extent of acetaldehyde adsorption as a function of contact time and temperature in the dark condition. It was observed that the amount of acetaldehyde adsorbed progressively increased with time and temperature. Equilibrium was achieved in 5 hr for a temperature higher than 5°C. In viewing all these kinetic curves, two phases were observed: (i) the first phase, where a rapid adsorption appeared within 1 hr contact time; (ii) the second phase, where a progressive adsorption occurred thereafter. In the fast adsorption phase, approximately 80% of acetaldehyde was removed. The subsequent slow adsorption stage was mainly caused by external active sites already being occupied by acetaldehyde, as well as by the slow diffusion of acetaldehyde into the pore spaces of N-TiO2. In general, adsorption involves a multistep process where the rate of adsorption is limited by intraparticle diffusion.Figure 3
(a) The fitting of pseudo-second-order kinetic plots for the adsorption of acetaldehyde onto N-TiO2 at various temperatures; the solid lines are the best fittings of PSO to the kinetic data. (b) Arrhenius equation plot for the adsorption of acetaldehyde onto N-TiO2. Experimental conditions: C
H
3
C
H
O
i = 176–1010 ppmv.The kinetic data was analyzed using PSO. The fittings of PSO to the kinetic data were presented as solid line in Figure3. Table 2 summarizes the amount of acetaldehyde adsorbed onto N-TiO2 sample at different temperatures and the corresponding fitting parameters of PSO model. Overall, the kinetic data correlate well with PSO models as judged from the high coefficient of determination values (R
2
>
0.95). The rate constant increased from 0.164 to 0.308 g/μmol/hr as temperature increased from 5 to 35°C. Thus, it was confirmed that an increase in temperature results in an increase in the adsorption rate. An increase in temperature may increase the driving force of diffusion across the external boundary layer and increase the rate of diffusion within the pores. Consequently, more acetaldehyde easily enters the interior pore of N-TiO2. The rate constants, k
2, listed in Table 2 can be related to temperature by Arrhenius equation and were used to determine the activation energy (E
a) of this adsorption process. From the Arrhenius plot of ln
⁡
k
2 versus 1
/
T (graph inset of Figure 3), an E
a of 14 kJ/mol was determined for the adsorption system. Normally a diffusion-controlled process has an E
a value lower than 42 kJ/mol [24]. The obtained E
a value is 14.0 kJ/mol, which indicates the adsorption of acetaldehyde that occurred during the diffusion control process [25].Table 2
Predicted constants of pseudo-second-order kinetic model for the adsorption of acetaldehyde onto N-TiO2 at various temperatures.
Temperature (°C)
Pseudo-second-order model
Activation energy
q
e (μmol/g)
k
2 (g/μmol/hr)
R
2
E
a (kJ/mol)
5
0.860
0.164
0.996
14.0
15
0.687
0.230
0.976
25
0.630
0.245
0.977
35
0.540
0.308
0.957


## 3.2.2. Adsorption Isotherms

Figure4 shows that the magnitude of acetaldehyde adsorption is proportional to temperature, and the results were fitted to the Langmuir isotherm model (equation (6)). Table 3 lists the Langmuir constants (i.e., q
m
a
x and K
L). The high coefficient of determination values confirms that the equilibrium data could be well represented by the Langmuir isotherm. The changes in the free energy (Δ
G°), enthalpy (Δ
H°), and entropy (Δ
S°) in the adsorption reaction system were calculated from the van’t Hoff plot (graph inset of Figure 4) and are listed in Table 3. The negative value of Δ
H° (−3.42 kJ/mol) indicates that the adsorption mechanism is exothermic in nature and that it was a diffusion-controlled process [26]. Moreover, the small value of Δ
H° implies weak bonding between acetaldehyde and the N-TiO2 surface. The negative value of Δ
G° and the positive value of Δ
S° (3.92 cal/K/mol) indicate that the adsorption of acetaldehyde onto N-TiO2 was a spontaneous process. The positive value of Δ
S° in this study suggests increased randomness at the gas-solid interface with some structural changes in the adsorbate and the adsorbent [27, 28]. In other words, a dissociative adsorption of acetaldehyde occurred possibly on the N-TiO2 surface during adsorption. But the low value of Δ
S° may imply that no remarkable change in entropy occurred during the adsorption of acetaldehyde onto N-TiO2. Normally, adsorption of gases causes a decrease in entropy owing to an orderly arrangement of the gas molecules on a surface. A slight increase in the Δ
G° values from −4.60 to −4.53 kcal/mol with a decrease in temperature from 35 to 5°C indicated that the adsorption process was more favorable and spontaneous at low temperatures.Table 3
Thermodynamic parameters for the adsorption of acetaldehyde onto N-TiO2 at various temperatures.
Temperature (°C)
q
m
a
x (μmol/g)
K
L (L/mol)
R
2
Δ
G
° (kcal/mol)
Δ
H
° (kcal/mol)
Δ
S
° (cal/K/mol)
5
196
0.0017
0.998
−4.53
−3.42
3.92
15
290
0.0018
0.979
−4.59
25
336
0.0022
0.992
−4.57
35
444
0.0036
0.982
−4.60Figure 4
(a) Adsorption isotherms of acetaldehyde adsorption on N-TiO2 at various temperatures. Solid lines are the best fit of Langmuir isotherm. (b) van’t Hoff plot for the adsorption of acetaldehyde onto N-TiO2.


## 3.3. Photooxidation Kinetics

### 3.3.1. Effect of Acetaldehyde Concentration

The influence of various acetaldehyde concentrations on the rate of acetaldehyde photooxidation is shown in Figure5, and the results correlated well with the L-H rate equation. Increasing the acetaldehyde concentration increases the concentration gradient between the bulk phase and the N-TiO2 surface where the adsorption of the reactant becomes the rate-controlling step. Thus, the higher the acetaldehyde concentration, the faster the reaction rate. According to the L-H model, the rate of the reactant molecular reaction is proportional to the fraction of surface covered by the reactants. Light intensity and temperature were held constant to determine the rate constant (k) (model I in Table 1). If the L-H form is valid, a plot of reciprocal of the degradation rate (1/rate) versus reciprocal of the initial acetaldehyde concentration will be linear. The graph inset of Figure 5 shows the corresponding plot of 1/rate versus 1/[CH3CHO] with good linearity. The red circles in Figure 5 represent a good fit of the experimental data to model I. It is clear that the L–H equation applied to one type of adsorption site which describes the initial acetaldehyde concentration dependencies satisfactorily. The values for k and K
a, generated by model fitting of the experimental data for a reaction temperature of 25°C, were 1.2666 μmol/cm2 and 0.0263 ppmv−1, respectively.Figure 5
The effect of acetaldehyde concentration on the photocatalytic reaction. Experimental conditions: temperature = 25°C,I
=
1.01 mW/cm2.


### 3.3.2. Effect of Visible-Light Intensity

The N-TiO2 absorbs light with a threshold wavelength that is enough to provide energy to overcome the band gap between the valence band and the conduction band. For N-TiO2, the wavelength of the visible light greater than 400 nm can provide enough energy to overcome the band gap (2.8 eV). At enough activation energy, the electrons will transfer between valence band and conduction band to form electron-hole pairs on the catalyst surface. The electron-hole pairs can trigger the photocatalytic oxidation reactions. Figure 6 depicts the effect of light intensity on the reaction rate of acetaldehyde. The effect of light intensity was examined under the reactant gas stream of 161 ppmv CH3CHO at 25°C. The light intensity used in this experiment was in the range of 0.4–3 mW/cm2. The reaction rates increased with increasing the light intensity, which can generate more photons and electron-hole pairs in the following series of free-radical reactions. The reaction rate constant is a function of light intensity according to a reported relationship [6]. By substituting the above relationship in model I under constant initial concentration of acetaldehyde, the rate expression should be(8)
r
=
k
′
′
I
α
k
′
′
=
k
′
K
a
C
a
1
+
K
a
C
a
.Figure 6
The effect of photon flux on the photocatalytic reaction. Experimental conditions: [CH3CHO] = 161 ppmv and temperature = 25°C.By plottinglog
⁡
(
r
) versus log
⁡
(
I
) (graph inset of Figure 6), a straight line with a slope (α) of 0.111 and intercept (log
⁡
k
′
′) of 0.1112 is obtained and then the rate constant (k
′) is 1.2918 μmol/cm2/hr/(mW/cm2)0.1112. The effect of light intensity on the photooxidation reaction was similar to that reported previously [29]. Photocatalytic reactors can be operated in two regimes with respect to light intensity: a first-order regime where the electron-hole pairs are consumed more rapidly by chemical reactions than by recombination, and a 0.5-order regime where the recombination rate dominates [30]. The present result was consistent with the above phenomenon indicating that the dominant mechanism is the electron-hole pair recombination reaction.To include explicit environmental condition dependence in the rate equation, the following assumptions were made: the Langmuir adsorption constants for monolayer adsorption on a homogeneous surface can be applied on predicting the photooxidation rate of the N-TiO2, and the rate constant (k) follows Arrhenius temperature dependence. With these assumptions, a photooxidation rate of the N-TiO2 is given by model II (Table 1). The constants of model II are summarized in Table 4 and the values calculated are in good agreement with the experimental data, as shown by blue squares in Figure 5. The model II was used to identify the range of two key parameters (i.e., acetaldehyde concentrations and visible-light intensity) that would lead to an enhanced acetaldehyde oxidation rate. As evident from Figure 7, when a photocatalytic device with the N-TiO2 catalyst operated at a maximum acetaldehyde concentration of approximately 250 ppmv, the oxidation rate would approach a constant value. Specifically, oxidation rate of the N-TiO2 photocatalyst increased with elevated acetaldehyde concentration and visible-light intensity.Table 4
Langmuir-Hinshelwood parameters for models I and II.
Parameter
Unit
Type of kinetic model
I
II
k
μmol/h
1.2666
—
K
a
1/ppmv
0.0263
0.0263
α
—
—
0.1110
k
′
μmol/cm2/hr/(mW/cm2)0.1112
—
1.2918Figure 7
Predicted oxidation rate dependence on acetaldehyde concentration and light intensity (model II).


## 3.3.1. Effect of Acetaldehyde Concentration

The influence of various acetaldehyde concentrations on the rate of acetaldehyde photooxidation is shown in Figure5, and the results correlated well with the L-H rate equation. Increasing the acetaldehyde concentration increases the concentration gradient between the bulk phase and the N-TiO2 surface where the adsorption of the reactant becomes the rate-controlling step. Thus, the higher the acetaldehyde concentration, the faster the reaction rate. According to the L-H model, the rate of the reactant molecular reaction is proportional to the fraction of surface covered by the reactants. Light intensity and temperature were held constant to determine the rate constant (k) (model I in Table 1). If the L-H form is valid, a plot of reciprocal of the degradation rate (1/rate) versus reciprocal of the initial acetaldehyde concentration will be linear. The graph inset of Figure 5 shows the corresponding plot of 1/rate versus 1/[CH3CHO] with good linearity. The red circles in Figure 5 represent a good fit of the experimental data to model I. It is clear that the L–H equation applied to one type of adsorption site which describes the initial acetaldehyde concentration dependencies satisfactorily. The values for k and K
a, generated by model fitting of the experimental data for a reaction temperature of 25°C, were 1.2666 μmol/cm2 and 0.0263 ppmv−1, respectively.Figure 5
The effect of acetaldehyde concentration on the photocatalytic reaction. Experimental conditions: temperature = 25°C,I
=
1.01 mW/cm2.


## 3.3.2. Effect of Visible-Light Intensity

The N-TiO2 absorbs light with a threshold wavelength that is enough to provide energy to overcome the band gap between the valence band and the conduction band. For N-TiO2, the wavelength of the visible light greater than 400 nm can provide enough energy to overcome the band gap (2.8 eV). At enough activation energy, the electrons will transfer between valence band and conduction band to form electron-hole pairs on the catalyst surface. The electron-hole pairs can trigger the photocatalytic oxidation reactions. Figure 6 depicts the effect of light intensity on the reaction rate of acetaldehyde. The effect of light intensity was examined under the reactant gas stream of 161 ppmv CH3CHO at 25°C. The light intensity used in this experiment was in the range of 0.4–3 mW/cm2. The reaction rates increased with increasing the light intensity, which can generate more photons and electron-hole pairs in the following series of free-radical reactions. The reaction rate constant is a function of light intensity according to a reported relationship [6]. By substituting the above relationship in model I under constant initial concentration of acetaldehyde, the rate expression should be(8)
r
=
k
′
′
I
α
k
′
′
=
k
′
K
a
C
a
1
+
K
a
C
a
.Figure 6
The effect of photon flux on the photocatalytic reaction. Experimental conditions: [CH3CHO] = 161 ppmv and temperature = 25°C.By plottinglog
⁡
(
r
) versus log
⁡
(
I
) (graph inset of Figure 6), a straight line with a slope (α) of 0.111 and intercept (log
⁡
k
′
′) of 0.1112 is obtained and then the rate constant (k
′) is 1.2918 μmol/cm2/hr/(mW/cm2)0.1112. The effect of light intensity on the photooxidation reaction was similar to that reported previously [29]. Photocatalytic reactors can be operated in two regimes with respect to light intensity: a first-order regime where the electron-hole pairs are consumed more rapidly by chemical reactions than by recombination, and a 0.5-order regime where the recombination rate dominates [30]. The present result was consistent with the above phenomenon indicating that the dominant mechanism is the electron-hole pair recombination reaction.To include explicit environmental condition dependence in the rate equation, the following assumptions were made: the Langmuir adsorption constants for monolayer adsorption on a homogeneous surface can be applied on predicting the photooxidation rate of the N-TiO2, and the rate constant (k) follows Arrhenius temperature dependence. With these assumptions, a photooxidation rate of the N-TiO2 is given by model II (Table 1). The constants of model II are summarized in Table 4 and the values calculated are in good agreement with the experimental data, as shown by blue squares in Figure 5. The model II was used to identify the range of two key parameters (i.e., acetaldehyde concentrations and visible-light intensity) that would lead to an enhanced acetaldehyde oxidation rate. As evident from Figure 7, when a photocatalytic device with the N-TiO2 catalyst operated at a maximum acetaldehyde concentration of approximately 250 ppmv, the oxidation rate would approach a constant value. Specifically, oxidation rate of the N-TiO2 photocatalyst increased with elevated acetaldehyde concentration and visible-light intensity.Table 4
Langmuir-Hinshelwood parameters for models I and II.
Parameter
Unit
Type of kinetic model
I
II
k
μmol/h
1.2666
—
K
a
1/ppmv
0.0263
0.0263
α
—
—
0.1110
k
′
μmol/cm2/hr/(mW/cm2)0.1112
—
1.2918Figure 7
Predicted oxidation rate dependence on acetaldehyde concentration and light intensity (model II).


## 3.4. Acetaldehyde Removal Mechanism in Indoor Environment Using the N-TiO2

To illustrate clearly the acetaldehyde removal mechanism of N-TiO2 under the indoor environment, a conceptual reaction scheme of the photocatalytic oxidation of the acetaldehyde using N-TiO2 nanomaterial is proposed (Figure 8). In the dark indoor environment, such as nighttime, the acetaldehyde pollutant will diffuse from air phase to the surface of N-TiO2 catalyst. The acetaldehyde would be absorbed onto N-TiO2 particles, and decreased temperature results in an increase in the adsorption rate. When the visible light was turned on, the N-TiO2 catalyst had shown a stronger absorption in the visible-light region of 400–500 nm wavelength possibly due to the fact that the band gap was narrowed by mixing of nitrogen 2p states with O 2p states on the top of the valence band or a creation of N-induced midgap level [31, 32]; the absorption at wavelength above 500 nm is ascribed to oxygen vacancy. The excited electron (e−) was generated from N-TiO2 and transferred to contact with acetaldehyde. At the same time, the remaining holes also migrate outward to form reactive species and react with acetaldehyde in indoor air. Therefore, the acetaldehyde was photooxidized by N-TiO2 and the photooxidation rate increased with elevated acetaldehyde concentration, but it was less sensitive to the changes of visible-light intensity.Figure 8
Conceptual reaction mechanism of N-TiO2 under an indoor environment for acetaldehyde removal.


## 4. Conclusions

Indoor acetaldehyde pollutant is harmful for human health. It is highly potential that the home appliances with visible-light-responsive TiO2 photocatalyst, such as N-TiO2, can be used to remove the acetaldehyde pollutant in future. In practical application, using the photocatalyst to remove the indoor acetaldehyde involves an absorption and photooxidation process in dark and visible-light irradiation, respectively. The key parameters of the removal of the acetaldehyde pollutant using N-TiO2 photocatalyst under the practical indoor environment conditions were investigated. The kinetics and mechanism of the removal of the acetaldehyde pollutant using N-TiO2 photocatalyst will be useful for applicability of these home appliances.An increase in temperature results in a decrease in the acetaldehyde adsorption rate. The adsorption reaction was a spontaneous and exothermic process. The reaction rate of acetaldehyde photooxidation increased with increasing acetaldehyde concentration and light intensity. The kinetic studies confirm that the reaction rate is successfully predicted using the L–H kinetic model. A reaction mechanism scheme for removal process of acetaldehyde using N-TiO2 under illumination of visible light was proposed. In conclusion, the prepared N-TiO2 photocatalyst has a high ability to absorb acetaldehyde in the dark conditions and has a significantly high photocatalytic activity under visible-light irradiation.



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*Source: 3058429-2016-04-28.xml*