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\DOI{10.5802/crchim.412}
\datereceived{2024-09-15}
\daterevised{2025-04-23}
\dateaccepted{2025-07-24}
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Adsorbent & 
\parbox[t]{11pc}{\centering $Q_{\max}$ (mg of nitrate/g) (original unit in the reference)} & 
\parbox[t]{6.5pc}{\centering $Q_{\max}$ (mg of nitrate nitrogen/g)}\vspace*{2pt} & 
Reference \\
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\dateposted{2025-10-02}
\begin{document}

\begin{noXML}

\CDRsetmeta{articletype}{research-article}

\title{Efficient nitrate removal from water by biosorption onto a
southern Mediterranean Sea native plant \textit{Posidonia oceanica}
(L.) powder}

\alttitle{\'{E}limination efficace des nitrates de l'eau par
biosorption sur une poudre de \textit{Posidonia oceanica} (L.), plante
indig\`{e}ne du sud de la M\'{e}diterran\'{e}e}

\author{\firstname{Abdelbasset} \lastname{Ben Dkhil}\CDRorcid{0009-0002-4704-1355}\IsCorresp}
\address{Laboratory of Water, Membrane and Environmental Biotechnology
LR15CERTE04, Center of Research and Water Technologies, University of
Carthage, Technopole of Borj Cedria, PB 273, 8020 Soliman, Tunisia}
\address{Department of Chemistry, Higher Institute of Sciences and
Technologies of Environment, University of Carthage, Borj Cedria, P.B.
no 1003, 2050 Hammam Lif, Tunisia}
\email[A. B. Dkhil]{alwaysoptimist20@gmail.com}

\author{\firstname{Haitham} \lastname{Elleuch}\CDRorcid{0000-0003-2373-6427}}
\address{Laboratory of Structural Organic Chemistry and Macromolecular
LR99ES14, Faculty of Sciences of Tunis, University of Tunis El Manar,
University Campus, 2092 Tunis, Tunisia}

\author{\firstname{Anissa} \lastname{Somrani}\CDRorcid{0000-0003-0193-2830}}
\address{Department of Physics, College of Science, Jouf University,
Sakaka, Aljouf 72341, Saudi Arabia}

\author{\firstname{Amor}\nobreakauthor\lastname{Hafiane}\CDRorcid{0000-0002-1198-4354}}
\addressSameAs{1}{Laboratory of Water, Membrane and Environmental
Biotechnology LR15CERTE04, Center of Research and Water Technologies,
University of Carthage, Technopole of Borj Cedria, PB 273, 8020
Soliman, Tunisia}

\shortrunauthors

\thanks{Ministry of Higher Education and Scientific Research of Tunisia}

\keywords{\kwd{Nitrate}\kwd{Biosorption}\kwd{\textit{P. oceanica}
(L.)}\kwd{Tunisian coasts}\kwd{Ecophysiology}\kwd{Isotherm
modeling}\kwd{Mechanism}}

\altkeywords{\kwd{Nitrates}\kwd{Biosorption}\kwd{\textit{Posidonia
oceanica} (L.)}\kwd{C\^{o}tes
tunisiennes}\kwd{\'{E}cophysiologie}\kwd{Mod\'{e}lisation
d'isotherme}\kwd{M\'{e}canisme}}

\begin{abstract}
An efficient method for nitrate removal from water was developed using
the sea plant \textit{Posidonia oceanica} (L.) as a biosorbent. Never
used previously for this purpose, this plant was applied in its natural
form without any chemical modification or cross-linking with another
material. The nitrate biosorption was carried out in batch experiment
at room temperature and under the optimum conditions of agitation time,
biosorbent dose, and pH. The Langmuir model fits the experimental
biosorption isotherm data better than the Freundlich one, confirming a
monolayer adsorption process onto homogeneous surface and giving a
relatively high maximum biosorption capacity 41.6 mg${\cdot}$g$^{-1}$
of nitrate nitrogen, thus demonstrating a good efficiency of the
biosorbent. The thermodynamic study shows not only that the biosorption
process is spontaneous and endothermic in nature, but also that
increasing temperature enhances the nitrate uptake. Furthermore, the
Dubinin--Radushkevich modeling gives a mean free energy of 15.81
kJ${\cdot}$mol$^{-1}$, indicating that nitrate was chemisorbed. The
same result was confirmed by kinetic modeling, which showed that the
pseudo-second-order model fits the experimental data better than the
pseudo-first order. The characterization by Fourier transform infrared
spectroscopy (FTIR), scanning electron microscopy (SEM), and
energy-dispersive X-ray spectroscopy (EDS) revealed that nitrate was
mainly retained on the biosorbent surface through electrostatic forces
with protonated amines  ($\mathrm{RNH}_{3}^{+}$;
$\mathrm{R}_{2}\mathrm{NH}_{2}^{+}$ and
$\mathrm{R}_{3}\mathrm{NH}^{+}$). However, ion exchange with hydroxide
ions and intraparticle diffusion may play an important role in the
biosorption process.
\end{abstract}

\begin{altabstract}
L'\'{e}limination efficace des nitrates de l'eau a \'{e}t\'{e}
r\'{e}alis\'{e}e par biosorption en utilisant la plante marine
\textit{Posidonia oceanica} (L.), qui n'avait jamais \'{e}t\'{e}
utilis\'{e}e auparavant pour cette application. Le nouveau biosorbant
de nitrate a \'{e}t\'{e} utilis\'{e} \`{a} l'\'{e}tat natif, sans
modification chimique ni r\'{e}ticulation avec un autre mat\'{e}riau.
La biosorption des nitrates a \'{e}t\'{e} effectu\'{e}e \`{a}
temp\'{e}rature ambiante, dans les conditions optimales de temps
d'agitation, de dose de biosorbant et du pH. Le mod\`{e}le de Langmuir
s'adapte mieux aux donn\'{e}es exp\'{e}rimentales de l'isotherme de
biosorption que celui de Freundlich, ce qui confirme un processus
d'adsorption monocouche sur une surface homog\`{e}ne. Il permet
\'{e}galement de d\'{e}terminer une capacit\'{e} de biosorption
maximale relativement \'{e}lev\'{e}e, atteignant 41,6
mg${\cdot}$g$^{-1}$ d'azote nitrique, montrant ainsi la bonne
efficacit\'{e} du biosorbant. L'\'{e}tude thermodynamique montre non
seulement que le processus d'adsorption est spontan\'{e}e et
endothermique mais aussi que l'augmentation de la temp\'{e}rature
am\'{e}liore l'absorption des nitrates. De plus, la mod\'{e}lisation
Dubinin--Radushkevich donne une \'{e}nergie libre moyenne de 15,81
kJ${\cdot}$mol$^{-1}$, ce qui signifie que les nitrates ont \'{e}t\'{e}
chimisorb\'{e}s. Le m\^{e}me r\'{e}sultat a \'{e}t\'{e} confirm\'{e}
par la mod\'{e}lisation cin\'{e}tique selon laquelle le mod\`{e}le de
pseudo-second ordre s'adapte mieux aux donn\'{e}es exp\'{e}rimentales
que celui de pseudo-premier ordre. La caract\'{e}risation par
spectroscopie infrarouge \`{a} transform\'{e}e de Fourier (FTIR),
microscopie \'{e}lectronique \`{a} balayage (MEB) et spectroscopie de
rayons X \`{a} dispersion d'\'{e}nergie (EDS) montre que les nitrates
ont \'{e}t\'{e} principalement retenus par {des} forces
\'{e}lectrostatiques avec les amines proton\'{e}es
($\mathrm{RNH}_{3}^{+}$ ; $\mathrm{R}_{2}\mathrm{NH}_{2}^{+}$ and
$\mathrm{R}_{3}\mathrm{NH}^{+}$) sur la surface biosorbante. Cependant,
l'\'{e}change d'ions avec les ions hydroxyde et la diffusion
intraparticulaire peuvent jouer un r\^{o}le important dans le processus
de biosorption. 
\end{altabstract}

\maketitle

{\pagebreak}

\twocolumngrid

\end{noXML}

\section{Introduction}\label{sec1}

The most common inorganic nitrogen species in the environment are
ammonia ($\mathrm{NH}_{3}$ or $\mathrm{NH}_{4}^{+}$), nitrite
$\mathrm{NO}_{2}^{-}$ and nitrate $\mathrm{NO}_{3}^{-}$. Ammonia is
always oxidized to nitrate through two successive steps by nitrifying
bacteria: $(\mathrm{NH}_{3}\text{ or } \mathrm{NH}_{4}^{+})\rightarrow
\mathrm{NO}_{2}^{-} \rightarrow  \mathrm{NO}_{3}^{-}$. As a result, the
nitrate concentration in water and soil is usually higher than that of
other inorganic nitrogen species~\cite{1}. Nitrate and other inorganic
nitrogen species are present in aquatic environments as normal
biological degradation products of proteins and nucleic acids. They can
also enter natural waters through agricultural fertilizers, animal
excrement runoff, industrial wastes, and sewage effluents~\cite{2,3}.
Due to its high solubility, nitrate is considered as one of the most
common contaminants in the ground water aquifers destined for drinking
water supplies~\cite{4}. In fact, the consumption of an excessive
amount of nitrate can lead to some diseases, among which the most
well-known is methemoglobinemia which affects bottle-fed infants under
six months (blue baby syndrome)~\cite{5,6}, and even children aged
between one and seven years~\cite{7}. The latter study which was
carried out in a North African country (Morocco) has shown that
blue-baby syndrome affects about 36.2\% of children who consume water
with nitrate concentrations above  50~mg/L~\cite{7}. Incidentally, a
study realized in Tunisia has revealed that groundwater in the north of
the country significantly exceeds this concentration
threshold~\cite{8}.\looseness=1 

Therefore, nitrate removal treatment of Tunisian groundwater is highly
recommended to protect the child population. Furthermore, nitrate can
be reduced to nitrite in acidic gastric medium, leading to the
formation of nitrosamine compounds, which are suspected to cause
gastric cancer among adult people~\cite{9}. A recent study conducted in
Tunisia between 2010 and 2014 has demonstrated that stomach cancer
rates are 61.4\% in men and 38.6\% in women. Among the most likely
direct causes is the consumption of foods rich in nitrites and
nitrates~\cite{10}. For this reason, the maximum admissible
concentration level of nitrate in drinking water should not exceed
10~mg${\cdot}$L$^{-1}$ of $\mathrm{NO}_{3}^{-}-\mathrm{N}$, as
recommended by the World Health Organization (WHO) and the U.S.
Environmental Protection Agency \mbox{\cite{11,12,13,14}.}\looseness=1 

Furthermore, the nitrate ions have a direct impact on the environment,
as a high input of nitrogen can cause a massive proliferation of
phytoplankton and algae in the aquatic environments, leading to
excessive consumption of dissolved oxygen. This phenomenon, called
eutrophication, causes the death and decay of many aquatic organisms,
causing in turn an imbalance of the ecosystem equilibrium~\cite{15,16}.
Currently, the main treatment methods used for nitrate removal from
natural water are reverse osmosis, electrodialysis, ion exchange,
biological processes, and adsorption. Not only have these methods
proven to be efficient, but they also provide high-quality water and
some of them are easy to apply. However, some are relatively expensive,
require high energy consumption and can be difficult to implement.
Moreover, some generate large volumes of wastewater, which presents an
additional challenge to address~\cite{17,18,19,20,21,22,23}. The
adsorption method is based on  the surface retention of nitrate ions
via ionic bonds or by intraparticle diffusion on a finely divided solid
material.

Compared to existing processes, adsorption is lower in cost, more
easily applied, cleaner, and safer without releasing nitrate-laden
wastewater or other chemicals. Currently, there is a wide range of
adsorbents used to remove nitrate from water which could be natural or
synthetic~\cite{24}. Natural adsorbents were used either in their
native state without any modification of surface structure such as
chitosan~\cite{25}, sugarcane bagasse~\cite{26}, agriculture
waste~\cite{27} and algae~\cite{28}, or thermally or chemically
activated such as activated carbon derived from pomegranate rind,
or~\cite{29} sepiolite activated by HCl~\cite{30}, etc. The synthetic
nitrate adsorbents can include polymer cross-linked with natural
materials~\cite{31}, activated materials coated with metal oxide
nanoparticles~\cite{32}, or chemical compounds cross-linked with a
polymer~\cite{33}. The maximum adsorption capacity of nitrate, obtained
with the adsorbents currently used, ranges from 0.25 to 350
mg${\cdot}$g$^{-1}$ of nitrate nitrogen according to the experimental
conditions in which the retention was carried out~\cite{24}. However,
the yield and effectiveness of an adsorbent can be better evaluated
objectively at the level of large-scale treatment station, where a
large amount of this material will be needed and high volume of water
will be treated. Consequently, the cost of the adsorbent will be a
leading factor which determines the economic profitability of the
denitrification process.\looseness=1 

Thus, due to the aforementioned reasons and within this context, the
present research work tries to investigate a natural low-cost adsorbent
for nitrate removal from aqueous solutions in batch experiments. The
adsorbent applied is a sea plant \textit{Posidonia oceanica} (L.)
widely abundant on Mediterranean coasts and available almost throughout
the year~\cite{34}. Apart from its nitrate biosorption properties in
its dead state, the living \textit{P.~oceanica} plays an important role
as a bioindicator for monitoring water quality, as well as for
assessing ecosystem health and evaluating the impacts of climate
change. In fact, \textit{P.~oceanica} meadows stabilize the sediment,
protect the beaches from erosion and have a positive effect on the
clarity of water~\cite{35,36,37}. As they are highly sensitive to
anthropogenic effects, the long-term decline in shoot density and the
morphological changes observed in \textit{P.~oceanica} meadows over
several decades, indicate climate change driven by the environmental
pollution~\cite{38,39,40,41}.

In the present work, the biosorbent was used in a dead state. It was
dried at room temperature, chipped and crushed into fine powder. The
method shows good efficiency, requires no chemicals except a few drops
of HCl or NaOH for pH adjustment, and does not release any waste
effluent, making it a promising alternative denitrification process
suitable for large-scale applications. Unlike reverse osmosis,
electrodialysis, ion exchange, and biological denitrification, the
adsorption process by \textit{P.~oceanica} requires neither
pretreatment of raw water except pH adjustment nor posttreatment of
water. This minimizes the treatment cost and increases the
competitivity of the process in comparison to other denitrification
processes cited previously. Furthermore, \textit{P.~oceanica} is
organic matter that can be incinerated after use and preserved in a
narrow place far from natural water resources, contributing to the
environmental protection.

\section{Material and methods}\label{sec2}
\subsection{Biosorbent preparation}\label{ssec21}

The seagrass \textit{P.~oceanica} was collected in the dead state as a
ball of rolled-up brown fibers from the Soliman delegation beach,
southeast of Tunis (GPS:{\unskip\break} Latitude: 30.732394, Longitude: 10.469454).
The harvest was done at the beginning of summer, in June. The
contribution of nitrates in this beach comes mainly from two sources:
the discharge of urban water treated by a plant that does not carry out
tertiary denitrification treatment and the runoff towards the sea of
the excrement of cows which are very numerous in this region. A mass of
100~g of sea plant was put in 500~mL of distilled water and stirred
during 24~h. Subsequently, the plant was dried at 80~\textdegree C,
chipped and crushed. The surface area of \textit{P.~oceanica} powder
was 25.75~m\tsup{2}${\cdot}$g$^{-1}$ according to the
Brunauer--Emmett--Teller (N\tsub{2}-BET) isotherm method and using an
Anton Pear autosorb 6100 analyzer. The powder obtained from the plant
was transferred into a concentrated NaOH solution (1~M) and the mixture
was stirred during 24~h~\cite{42,43,44,45}. The purpose of this
operation is to force an ionic exchange between hydroxide ions
($\mathrm{OH}^{-}$) and the nitrate ions ($\mathrm{NO}_{3}^{-}$)
already accumulated by the \textit{Posidonia} fibers when it was in the
sea. This ionic exchange could also clear \textit{P.~oceanica} of other
anions such as:  $\mathrm{Cl}^{-}$, $\mathrm{F}^{-}$, $\mathrm{I}^{-}$,
$\mathrm{SO}_{4}^{-}$, etc., some of which disturb the determination of
nitrate by the sodium salicylate method~\cite{46} and compete with it
on the active sites of the biosorbent. Eventually, the biosorbent
powder was washed with deionized water in excess, dried at
40~\textdegree C in an oven, and preserved in a desiccator for further
use.

\subsection{Reagents}\label{ssec22}

All chemicals used were of analytical grade and all nitrate
concentrations were expressed in nitrate nitrogen, such as
1~mg${\cdot}$L$^{-1}$ of $\mathrm{NO}_{3}^{-}-\mathrm{N}$ corresponds to
4.428~mg${\cdot}$L$^{-1}$ of $\mathrm{NO}_{3}^{-}$~\cite{46}. The nitrate
stock solution having a concentration of 100~mg${\cdot}$L$^{-1}$ of 
$\mathrm{NO}_{3}^{-}-\mathrm{N}$ was prepared by dissolving 0.772~g of 
KNO\tsub{3} previously dried at 105~\textdegree C for 4~h, in pure
water  adjusted to a 1000~mL volumetric flask~\cite{1}. The stock
solution was stored in a refrigerator at 4~\textdegree C and prepared
anew every two months. Working standards ranging from 0.5 to
5~mg${\cdot}$L$^{-1}$ of $\mathrm{NO}_{3}^{-}-\mathrm{N}$ were prepared
immediately before usage by accurate serial dilution of the stock
solution. The other solutions required for the nitrate calibration
curve or to carry out the adsorption experiments were prepared by
dissolving the corresponding salt in the appropriate volume of
deionized water. The pH was adjusted using HCl or NaOH (0.1~M)
solutions.

\subsection{Nitrate determination}\label{ssec23}

Nitrate concentration after each biosorption experiment was determined
by the sodium salicylate spectrophotometric method described by many
previous works~\cite{46,47,48}. Absorbance measurements were performed
using a double beam Shimadzu 1800 UV--Vis spectrophotometer. Direct
absorbance at 415~nm was recorded for each nitrate solution using a
pair of UV quartz cells with 1~cm in optical path length. Nitrate
concentration was deduced by the projection on a calibration curve
plotting absorbance versus concentration. An example of such a
calibration curve is given by Figure~\ref{fig1}.

\begin{figure}
\includegraphics{fig01}
\caption{\label{fig1}Calibration curve plotting absorbance versus
concentration of standard solutions.}
\end{figure}

The linear fit of the points in Figure~\ref{fig1} gives a slope of
0.28958, an intercept of ${-}$0.00584 and a correlation coefficient of
0.9998. The latter parameter presents a strong linear correlation
between absorbances and concentrations, confirming the accuracy of the
nitrate determination method used.

\subsection{Batch biosorption experiments}\label{ssec24}

Batch experiment is an analytical method conducted in static mode,
which involves contacting a fixed volume of water to purify in contact
with a given mass of biosorbent, under controlled conditions of
agitation time, temperature, and pH. For all experiments, the biomass
dosage of \textit{P.~oceanica} was chosen at its optimum value and
biosorption was conducted at room temperature (30 ${\pm}$ 2~\textdegree
C), in 100~mL erlenmeyer flasks and under 100~rpm agitation. A given
mass of \textit{P.~oceanica} was agitated in 30~mL of a nitrate
solution using an LBX SO4X10 multiposition magnetic stirrer. The
solution's pH was measured by an AZ 86505 pH meter and adjusted using
HCl or NaOH (0.1~M) solutions as required. The zeta potential was
determined using a Malvern S90 zetasizer. 

\subsection{Biosorption isotherms}\label{ssec25} 

Three isotherm models (Langmuir, Freundlich, and Dubinin--Radushkevich)
were investigated to fit the sorption equilibrium data obtained in this
research work. The comparison of the correlation coefficient ($r$) of
the linear forms of all models indicates which isotherm provides the
best fit of the experimental data and therefore best describes the
sorption process. 

The Langmuir isotherm describes quantitatively the formation of a
monolayer adsorbate on the outer surface of the adsorbent, and no
subsequent further adsorption takes place. This model assumes monolayer
adsorption onto a homogeneous surface with a finite number of identical
sites, and without interaction between sorbed molecules. Based on these
assumptions, Langmuir established the equation of his model as follows
(Equation~(\ref{eq1}))~\cite{49,50}:
{\begin{equation}\label{eq1}
Q_{\mathrm{e}} = \dfrac{Q_{\max}bC_{\mathrm{e}}}{1 + bC_{\mathrm{e}}},
\end{equation}}\unskip
where $Q_{\mathrm{e}}$ (mg${\cdot}$g$^{-1}$) is the amount of solute 
adsorbed by adsorbent at equilibrium, $Q_{\max}$ (mg${\cdot}$g$^{-1}$) 
is the maximum adsorption capacity, $C_{\mathrm{e}}$
(mg${\cdot}$L$^{-1}$)  is the solute concentration at equilibrium and
$b$ (L${\cdot}$mg$^{-1}$) is the Langmuir constant related to the
adsorption energy. The linear  form is represented by the following
equation:
{\begin{equation}\label{eq2}
\dfrac{1}{Q_{\mathrm{e}}} = \dfrac{1}{Q_{\max}} + 
\dfrac{1}{bQ_{\max}C_{\mathrm{e}}}.
\end{equation}}\unskip
The essential feature of the Langmuir isotherm may be expressed as a
dimensionless equilibrium factor $R_{\mathrm{L}}$, calculated as
follows:
{\begin{equation}\label{eq3}
R_{\mathrm{L}} = \dfrac{1}{1+bC_{0}}.
\end{equation}}\unskip
The value of $R_{\mathrm{L}}$ indicates the adsorption nature to be
unfavorable ($R_{\mathrm{L}} > 1$), linear ($R_{\mathrm{L}} = 1$),
favorable  ($0 < R_{\mathrm{L}} < 1$), or irreversible 
($R_{\mathrm{L}} = 0$)~\cite{51}.
The Freundlich isotherm is an empirical model based on the adsorption
on a heterogeneous surface and the adsorbate possibly multilayered. It
is expressed by the following equation~\cite{50,52}:
{\begin{equation}\label{eq4}
Q_{\mathrm{e}} = K_{\mathrm{F}}{C}_{\mathrm{e}}^{1/n},
\end{equation}}\unskip
where $K_{\mathrm{F}}$ is the Freundlich constant indicative of
adsorption capacity and $n$ is a factor measuring the degree of
heterogeneity of adsorption sites. The value of $n$ indicates whether
the adsorption is good ($2<n<10$), moderately  difficult ($1<n<2$), or
poor ($n<1$)~\cite{53}. The linear form of this model is given as
follows:
{\begin{equation}\label{eq5}
\mathrm{Ln}(Q_{\mathrm{e}}) = \mathrm{Ln} (K_{\mathrm{F}}) +
\dfrac{1}{n} \mathrm{Ln} (C_{\mathrm{e}}).
\end{equation}}\unskip
The Dubinin--Radushkevich (D--R) isotherm is generally applied to
determine whether the nature of adsorption is physical or chemical. The
D--R isotherm is more general than the Langmuir isotherm, as it is not
based on ideal assumptions such as equipotential of the adsorption
sites, the absence of steric hindrance between adsorbates and incoming
particles, and surface homogeneity on a microscopic level~\cite{54,55}.
The D--R isotherm gives insights into the adsorption mechanism based on
potential theory and assuming a heterogeneous surface of the solid
phase~\cite{55,56}. The linear form of the D--R isotherm is described
by the following equation~\cite{57,58}:
{\begin{equation}\label{eq6}
\mathrm{Ln}(Q_{\mathrm{e}}) = \mathrm{Ln}(Q_{\max}) - \beta
\varepsilon^{2},
\end{equation}}\unskip
where $Q_{\mathrm{e}}$ (mol${\cdot}$g$^{-1}$) is the amount of solute
retained by adsorbent at equilibrium, $Q_{\max}$ (mol${\cdot}$g$^{-1}$)
the theoretical saturation capacity, $\beta$
(mol\tsup{2}${\cdot}$J$^{-2}$)  an activity coefficient related to the
mean free energy of adsorption and  $\varepsilon$
(J${\cdot}$mol$^{-1}$) is the Polanyi potential defined for  the
liquid/solid adsorption as follows~\cite{59}:
{\begin{equation}\label{eq7}
\varepsilon = RT \mathrm{Ln} \left(1 + \dfrac{1}{C_{\mathrm{e}}}\right),
\end{equation}}\unskip
where $R$ (8.314~J${\cdot}$mol$^{-1}{\cdot}$K$^{-1}$) is the
universal gas constant, $T$ (K) the absolute temperature and
$C_{\mathrm{e}}$ (mol${\cdot}$L$^{-1}$) the concentration of solute at
equilibrium. The essential feature of the D--R isotherm is the mean
free energy $E$ (J${\cdot}$mol$^{-1}$), computed as follows:
{\begin{equation}\label{eq8}
E = \dfrac{1}{\sqrt{2\beta}}.
\end{equation}}\unskip
The value of $E$ indicates whether the adsorption mechanism occurs as
physisorption ($E<8$~kJ${\cdot}$mol$^{-1}$) or chemisorption ($8<E<16$
kJ${\cdot}$mol$^{-1}$)~\cite{59,60}. Consistently with
Equation~(\ref{eq6}), the plot Ln($Q_{\mathrm{e}}$) versus
$\varepsilon^{2}$ allows the determination of $\beta$ and $Q_{\max}$
through the slope and the intercept, respectively.

\subsection{Chi-square test}\label{ssec26}

The statistical hypothesis chi-square test ($\rmchi^{2}$) is applied in
this work to verify the conformity between the experimental data and
the values predicted by a theoretical model. The comparison is made
between the experimental equilibrium biosorption capacities and those
calculated by the model.

The conformity chi-square test is based on two hypotheses, as follows:

Hypothesis $\mathrm{H}_{0}$: There is conformity with the theoretical
model at a given confidence level.

Hypothesis $\mathrm{H}_{1}$: There is no conformity with the
theoretical model.

The observed chi-square value is calculated as follows:
{\vspace*{4pt}}
{\begin{equation}\label{eq9}
\rmchi _{\mathrm{obs}}^{2} = \displaystyle\sum _{i=1}^{n}
\dfrac{(Q_{\mathrm{e}(\exp)} - Q_{\mathrm{e}(\mathrm{theo})})^{2}}
{Q_{\mathrm{e}(\mathrm{theo})}},
\end{equation}
{\vspace*{4pt}}}\unskip
where $Q_{\mathrm{e}(\exp)}$ and $Q_{\mathrm{e}(\mathrm{theo})}$ are
the experimental and calculated equilibrium biosorption capacities,
respectively, $n$ the number of measures, and  ($n-1$) the number of
degrees of freedom.

The comparison between the observed chi-square value 
$\rmchi_{\mathrm{obs}}^{2}$ and the critical chi-square value 
$\rmchi_{\mathrm{c}}^{2}$, at a given number of degrees of freedom and
confidence level, can lead to two possibilities:
{\begin{eqnarray*}
\rmchi _{\mathrm{obs}}^{2} &\leq&  \rmchi _{\mathrm{c}}^{2}\colon
\mbox{ hypothesis }\mathrm{H}_{0} \mbox{ is retained} \\
\rmchi _{\mathrm{obs}}^{2} &>& \rmchi _{\mathrm{c}}^{2}\colon
\mbox{ hypothesis }\mathrm{H}_{1} \mbox{ is retained}. 
\end{eqnarray*}}\unskip

\subsection{Biosorption kinetic modeling}\label{ssec27}

In this research work, we applied the well-known and much-cited
following kinetic models: pseudo-first order~\cite{61}, pseudo-second
order~\cite{62}, and intraparticle diffusion models~\cite{63}.

Table~\ref{tab1} shows the equation of each kinetic model and its
integrated form, where $Q_{{t}}$ and $Q_{\mathrm{e}}$ 
(mg${\cdot}$g$^{-1}$) are the amount of solute at an instant $t$ (min)
and at equilibrium, respectively, $C_{\mathrm{i}}$ 
(mg${\cdot}$g$^{-1}$) is a constant that depends on adsorbent nature, 
$k_{1}$ (min$^{-1}$), $k_{2}$
(g${\cdot}$mg$^{-1}{\cdot}$min$^{-1}$), and $k_{\mathrm{p}}$
(mg${\cdot}$g$^{-1}{\cdot}$min$^{-0.5}$) are the rate constant of the
pseudo-first order, pseudo-second order, and intraparticle diffusion
models, respectively.  $Q_{\mathrm{cal.1}}$ and $Q_{\mathrm{cal.2}}$
are the calculated (theoretical) amounts of solute retained at
adsorption equilibrium for the pseudo-first order and pseudo-second
order models, respectively. The coefficients  $k_{1}$ and
$Q_{\mathrm{cal.1}}$, $k_{2}$ and $Q_{\mathrm{cal.2}}$, and
$k_{\mathrm{p}}$ can be estimated from the curve plotting
Ln($Q_{\mathrm{e}} - Q_{{t}}$) versus $t$, $t/Q_{{t}}$
versus $t$, and $Q_{{t}}$ versus $t^{0.5}$ for the pseudo-first
order, pseudo-second order, and intraparticle diffusion models,
respectively (Table~\ref{tab1}).

%tab1
\begin{table*}
\caption{\label{tab1}Original, linearized forms and deduced parameters
of the kinetic models used\vspace*{2pt}}
\tabcolsep=3.5pt\fontsize{9.6}{12}\selectfont
\begin{tabular}{ccccc}
\thead
Kinetic model & Original form & Linearized form & Plot & 
\parbox[t]{5.5pc}{\centering Deduced parameters}\vspace*{2pt} \\
\endthead
\parbox[t]{5.5pc}{\centering Pseudo-first order model} & 
$Q_{t} = Q_{\mathrm{e}}(1 - \exp (-k_{1}t))$ &
$\mathrm{Ln}(Q_{\mathrm{e}}-Q_{{t}}) = 
\mathrm{Ln}(Q_{\mathrm{cal}.1})-k_{1}t$ & 
$\mathrm{Ln}(Q_{\mathrm{e}}-Q_{{t}})$ versus $t$ & 
\parbox[t]{5.5pc}{\centering Slope: $k_{1}$\Lbreak Intercept: $Q_{\mathrm{cal}.1}$}  \\

\parbox[t]{5.5pc}{\centering Pseudo-second order model} & 
$Q_{{t}} = \dfrac{k_{2}{Q}_{t}^{2}t}{1+k_{2}Q_{\mathrm{e}}t}$ &
$\dfrac{t}{Q_{t}} = \dfrac{1}{k_{2}{Q}_{\mathrm{cal}.2}^{2}} + 
\dfrac{1}{Q_{\mathrm{cal}.2}}t$ & 
$\dfrac{t}{Q_{{t}}}$ versus $t$ & 
\parbox[t]{5.5pc}{\centering Slope: $k_{2}$\Lbreak Intercept: $Q_{\mathrm{cal}.2}$}\vspace*{4pt} \\

\parbox[t]{5.5pc}{\centering Intraparticle diffusion model} &
$Q_{{t}} = k_{\mathrm{p}}\sqrt{t} + C_{\mathrm{I}}$ &
$Q_{t} = k_{\mathrm{p}}\sqrt{t} + C_{\mathrm{I}}$ & 
$Q_{t}$ versus $t$ & Slope: $k_{\mathrm{p}}$
\vspace*{2pt}
\botline
\end{tabular}
\vspace*{8pt}
\end{table*}

Among the first two models, the best fit was selected based on both
linear regression correlation coefficient ($r$) of the plot and the
concordance between the calculated and experimental equilibrium amount
values ($Q_{\mathrm{e}}$) of adsorbed solute. For the third model, the
shape of the curve $Q_{{t}}$ versus $t^{0.5}$ determines the
importance of each transport step for solute adsorption. 

\subsection{Thermodynamic parameters}\label{ssec28}

The thermodynamic parameters of the biosorption reaction such as
equilibrium constant $K_{\mathrm{C}}$, standard free enthalpy change 
$\Delta {G}_{T}^{0}$, standard enthalpy $\Delta {H}_{T}^{0}$
change and standard entropy change $\Delta {S}_{T}^{0}$ were
determined using the relations summarized in Table~\ref{tab2}.

%tab2
\begin{table*}
\caption{\label{tab2}Thermodynamic relations used\vspace*{2pt}}
\begin{tabular}{lc}
\tbody
Equilibrium biosorption constant & 
$K_{\mathrm{C}} = \dfrac{(C_{0}-C_{\mathrm{e}})}{C_{\mathrm{e}}} = 
\dfrac{C_{\mathrm{S}}}{C_{\mathrm{e}}}$\raisebox{15pt}{} \\

Standard free enthalpy change & 
$\Delta {G}_{T}^{0} = -RT\mathrm{Ln}(K_{\mathrm{C}}) = 
\Delta {H}_{T}^{0} - T\cdot\Delta {S}_{T}^{0}$\vspace*{3pt} \\

Van't Hoff relation & 
$\mathrm{Ln}(K_{\mathrm{C}}) = \dfrac{\Delta {S}_{T}^{0}}{R} - 
\dfrac{\Delta {H}_{T}^{0}}{RT}$
\vspace*{2pt}
\botline
\end{tabular}
\xxtabnote{Where $C_{\mathrm{e}}$ is the
concentration of $\mathrm{NO}_{3}^{-}-\mathrm{N}$ (mg${\cdot}$L$^{-1}$)
at equilibrium and $C_{\mathrm{S}}$ the adsorbed concentration of
$\mathrm{NO}_{3}^{-}-\mathrm{N}$  (mg${\cdot}$L$^{-1}$) by biosorbent.
The concentration $C_{\mathrm{S}}$ also represents the mass of
adsorbate retained by adsorbent mass corresponding to one liter of
solution~\cite{64}.\vspace*{2pt}}
\end{table*}

\subsection{Infrared spectroscopy analysis}\label{ssec29}

Infrared spectroscopy was used to determine the main functional groups
in \textit{P.~oceanica} powder involved in the biosorption process. The
powder of \textit{P.~oceanica} (3~mg) was dispersed in KBr (200~mg,
spectroscopic grade), and the mixture was ground in an agate mortar
until perfectly homogeneous. The mixture was transferred to a disc of
diameter 1~cm under a pressure of 14 tons using a hydraulic press. The
infrared spectra were recorded by scanning the range of 
400--4000~cm$^{-1}$ using a PerkinElmer Frontier FTIR spectrometer.
The reference spectrum was a polystyrene film and the background
spectrum of pure KBr was subtracted from the sample spectrum. Each
recorded spectrum was the result of a ten-scan accumulation and was
obtained with a 1~cm$^{-1}$ spectral resolution. 

Scheme~\ref{sch1} shows the two vibrational modes of nitrates that
appear in the mid-infrared spectral region, whose corresponding bands
could be very helpful in the characterization.

\begin{scheme}
\includegraphics{sc01}
\vspace*{7pt}
\caption{\label{sch1}Valence vibration of the nitro --NO\tsub{2} group.}
\end{scheme}

\subsection{Analysis by scanning electron microscopy and energy
dispersive spectroscopy}\label{ssec210}

The analysis by scanning electron microscopy (SEM) gives information
about the topography change on the \textit{P.~oceanica} surface before
and after nitrate biosorption. Likewise, the analysis by
energy-dispersive X-ray spectroscopy (EDS) detects the slightest
variation of chemical elemental composition of biosorbent surface. The
two types of analysis were achieved by a ThermoFisher electron
microscope at 100 and 10~$\upmu$m resolution level, with 5~kV
acceleration voltage, 500~$\upmu$m horizontal field width and under a
vacuum of 10$^{-3}$~Pa.

\section{Results and discussion}\label{sec3}
\subsection{Effect of experimental conditions}\label{ssec31}
\subsubsection{Effect of solution pH}\label{sssec311}

The initial pH is a crucial factor that should be taken into
consideration  in biosorption processes, as it provides information
about the protonation  state of active sites on the biosorbent surface
and the possible influence  of free hydronium ions (H\tsub{3}O$^{+}$)
and free hydroxide ions  (OH$^{-}$) on the retention of nitrates.
Firstly, the pH effect on zeta  potential of water-suspended biosorbent
is worth exploring to gain insight  into its surface charge for each pH
range. The zeta potential measurements  of aqueous solutions containing
suspended \textit{P.~oceanica} powder are  displayed against pH in
Figure~\ref{fig2}. In pH ranging from 2 to 6, the  zeta potential is
positive with gradual decrease, thus confirming the presence of
positively charged groups ($\mathrm{RNH}_{3}^{+}$, 
$\mathrm{R}_{2}\mathrm{NH}_{2}^{+}$,  $\mathrm{R}_{3}\mathrm{NH}^{+}$,
and  $-\mathrm{OH}_{2}^{+}$) on the biosorbent surface. The more
neutralized  the protonated groups are, the more the solution pH
increases. Indeed, the rapid decrease in the zeta potential between pH
${=}$\ 6 and 6.5 can be attributed to the significant loss of positively
charged groups neutralized  by hydroxide ions. The zero-crossing point
of the plot coincides with the pH value of 6.5. Commonly known as
zero-charge point, this is when the  surface of the biosorbent
suspended in water is perfectly neutral, due to compensation between
negative and positive charges on its surface. For pH values higher than
the zero-charge point, the adsorbent surface becomes  negatively
charged due to the accumulation of hydroxide ions, which may cause the
appearance of conjugate bases for some organic groups.

\begin{figure}
\vspace*{-4pt}
\includegraphics{fig02}
\vspace*{-4pt}
\caption{\label{fig2}Effect of pH on zeta potential of
\textit{P.~oceanica} powder in aqueous suspension. (Biomass dose ${=}$\
 3.33~g${\cdot}$L$^{-1}$, temperature ${=}$\ 30 ${\pm}$ 2~\textdegree C).}
\vspace*{-4pt}
\end{figure}

The pH effect on the biosorbed amount was investigated over the pH
range from 2 to 12 and results are presented in Figure~\ref{fig3}.

\begin{figure}
\vspace*{-4pt}
\includegraphics{fig03}
\vspace*{-4pt}
\caption{\label{fig3}Effect of pH on the nitrate biosorption onto
\textit{P.~oceanica} powder. (Initial aqueous concentration ${=}$\ 
5~mg${\cdot}$L$^{-1}$, Biomass dose ${=}$\ 3.33~g${\cdot}$L$^{-1}$,
contact time ${=}$\ 60 min, temperature ${=}$\ 30 ${\pm}$ 2~\textdegree
C).}
\vspace*{-4pt}
\end{figure}

Three significant regions characterize the plot. The first one covers
the pH acid range from 2 to 5 where the biosorption percentage tends to
be constant with a slight variation between pH ${=}$\ 2 and 3. In this
region, free H\tsub{3}O$^{+}$ react with superficial functional
groups of \textit{P.~oceanica} such as amine and alcohol and protonate
them affording ionized functions $\mathrm{RNH}_{3}^{+}$,
$\mathrm{R}_{2}\mathrm{NH}_{2}^{+}$, $\mathrm{R}_{3}\mathrm{NH}^{+}$,
and $-\mathrm{OH}_{2}^{+}$, which retain \mbox{nitrate}
($\mathrm{NO}_{3}^{-}$) by electrostatic forces. This result is in good
agreement with the one obtained from the zeta potential study. However,
at very low pH values between 3 and 2, the slight decrease of
biosorption percentage could be explained by the increase in the number
of hydrogen bonds between nitrate and hydronium ions in
solution~\cite{65}. It can also be ascribed to the partial degradation
of cellulose and other polymeric chains in the biosorbent~\cite{66}.

The second region covers the pH range from 5 to 7 and presents a great
increase in biosorption percentage, reaching 72.3\% at pH ${=}$\ 6. The
ascending part of the curve could be explained by a cooperative effect
between two retention modes of nitrate happening simultaneously. The
electrostatic interaction with ionized functions and ion exchange with
hydroxide ions (OH$^{-}$) retained by \textit{P.~oceanica}  while in
seawater and likely during the Section~\ref{ssec21} step of
biosorbent preparation. The cooperative effect between ion exchange and
electrostatic interactions to retain nitrate is also confirmed by the
zeta potential study. Indeed, at pH ${=}$\ 6, the adsorbent surface is
positively charged, and the ion exchange between hydroxide and nitrate
can easily take place. The descending part of the curve could be
explained by the progressive disappearance of ionized sites, especially
$-\mathrm{NH}_{3}^{+}$ which decreases by about 90\% at pH ${=}$\ 7 as
described by Navarro et~al.~\cite{67}.

The third region covers the pH range from 7 to 12, showing a nearly
constant biosorption percentage with a very small increase toward the
most basic media. In this pH region, the nitrate retention is mainly
carried out by ion exchange with hydroxide ions. The ion exchange with
hydroxide ions was experimentally proven by an increase in the final pH
of solutions with initial pH values ranging from 5 to 12. As a result
of this study, the optimum pH value chosen was 6, which will be used
for all further experiments.

\begin{figure}
\includegraphics{fig04}
\caption{\label{fig4}Effect of the biomass dosage on the nitrate
biosorption onto \textit{P.~oceanica} powder. (Initial aqueous
concentration ${=}$\ 5~mg${\cdot}$L$^{-1}$, pH ${=}$\ 6, contact time ${=}$\
 60 min, temperature ${=}$\ 30 ${\pm}$ 2~\textdegree C).}
\end{figure}

\subsubsection{Effect of biosorbent dosage}\label{sssec312}

The effect of \textit{P.~oceanica} powder dose on the biosorption of
nitrate is displayed in Figure~\ref{fig4}. The maximum biosorption
percentage (82.2\%) was found for the lowest biomass dosage (0.167
g${\cdot}$L$^{-1}$), and subsequently, the biosorption significantly
decreases until reaching 65\% at 6 g${\cdot}$L$^{-1}$. It is worth
noting that the biosorbent dose used arbitrarily in the study of the pH
effects (3.333 g${\cdot}$L$^{-1}$) corresponds to 73.0\%, which is not
the optimum value. Figure~\ref{fig4} reveals that this value is located
midway between the highest and lowest biosorption percentage values. On
the other hand, the decrease in biosorption percentage versus the
increasing biosorbent dose, could be explained by the fact that higher
concentrations of biomass in suspension can exert a shell effect, thus
preventing a good contact between nitrate and available adsorption
sites. In fact, the shell effect can be understood by an agglomeration
of adsorbent grains on top of each other when they are dispersed at
high density in the solution, leading to a remarkable reduction in
their specific surface area. As a result, there will not be good
contact between the nitrate ions and the active sites of the
biosorbent, which leads to a decrease in the percentage of biosorption
when the biomass dose exceeds certain \mbox{limits.} Hence, biosorbent powder
exhibits greater nitrate retention at lower dispersion densities than
at higher ones. Similar findings were reported and interpreted in
previous research works, such as ammonium biosorption by
\textit{P.~oceanica}~\cite{68}, orthophosphate biosorption by date palm
fibers~\cite{69}, and heavy metals biosorption by algae~\cite{70,71}.
As a result of this study, the optimum biosorbent dose was found to be
0.167 g${\cdot}$L$^{-1}$, which would be used for all further\break
experiments. 

\subsubsection{Effect of initial concentrations and contact time}
\label{sssec313}

The effect of initial concentration and contact time on the biosorption
of nitrate is displayed in Figure~\ref{fig5}. The uptake efficiency is
expressed by the amount of nitrate nitrogen biosorbed
($Q_{\mathrm{e}}$). The experiments were carried out at optimum
conditions of pH and biomass dose and at room temperature. All plots in
Figure~\ref{fig5} show almost the same trend with three kinetic stages.
The first one corresponds to the first 10 min with a rapid increase in
the biosorbed amount. In this stage, the fast biosorption can be
explained by the existence of a large amount of sorption sites that
provide the driving force for faster transfer of the nitrate onto the
biosorbent surface as described in a previous work using active carbon
to adsorb lead~\cite{72}. The second stage covers the range from 10 to
60 min and presents a slow and gradual increase in the biosorbed
amount. This can be ascribed to the progressive saturation of sorption
sites and the reaching of the equilibrium at the end of this
period~\cite{72}. The third stage starts from 60 min and presents a
saturation plateau corresponding to biosorption equilibrium.
Furthermore, raising the nitrate concentration from 5 to 14
mg${\cdot}$L$^{-1}$ of $\mathrm{NO}_{3}^{-}-\mathrm{N}$, allows the 
\textit{P.~oceanica} powder to increase its biosorption capacity.

\begin{figure}
\includegraphics{fig05}
{\vspace*{-2pt}}
\caption{\label{fig5}Effect of initial concentrations and contact time
on the nitrate biosorption onto \textit{P.~oceanica} powder. (pH ${=}$\
 6, Biomass dose ${=}$\ 0.167 g${\cdot}$L$^{-1}$, temperature ${=}$\ 30
${\pm}$ 2~\textdegree C).}
{\vspace*{-2pt}}
\end{figure}

Similar findings of the kinetic stages and the effect of initial
concentrations were reported for the biosorption of chromium by
\textit{P.~oceanica} fibers~\cite{42}. Thus, the contact time 60 min
represents the beginning of the equilibrium state. As a result of this
study, the efficiency of biosorption increases with the increase in the
initial nitrate concentration, and the optimum contact time is 60 min.
It is noted that the optimal contact time coincides with the agitation
time chosen arbitrarily in previous experiments (effect of pH and
biosorbent dosage), and this time would be used in the following
experiments.

\subsubsection{Effect of temperature}\label{sssec314}

The effect of the temperature on the biosorption of nitrate by
\textit{P.~oceanica} powder is displayed in Figure~\ref{fig6}. This
figure shows that the biosorption \mbox{percentage} increases
significantly with the increase in temperature, from 82\% at
30~\textdegree C to 87\% at 50~\textdegree C. Such a behavior means
that biosorption is an endothermic process and any increase in
temperature is favorable to nitrate retention by \textit{P.~oceanica}
powder.

\begin{figure}
{\vspace*{-2pt}}
\includegraphics{fig06}
\caption{\label{fig6}Effect of temperature on the nitrate biosorption
onto \textit{P.~oceanica} powder. (Initial aqueous concentration ${=}$\
 5~mg${\cdot}$L$^{-1}$, pH ${=}$\ 6, Biomass dose ${=}$\ 0.167
g${\cdot}$L$^{-1}$, contact time ${=}$\ 60~min).}
{\vspace*{-6pt}}
\end{figure}

\subsection{Biosorption isotherms}\label{ssec32}

The biosorption isotherm was plotted under the optimal conditions of
different factors, at room \mbox{temperature} and with initial nitrate
concentrations ranging from 5 to 15 mg${\cdot}$L$^{-1}$ of 
$\mathrm{NO}_{3}^{-}-\mathrm{N}$ with an increment of 0.5
mg${\cdot}$L$^{-1}$ of $\mathrm{NO}_{3}^{-}-\mathrm{N}$, as presented in
Figure~\ref{fig7}.  The Langmuir linear model applied to this isotherm
plotting $1/Q_{\mathrm{e}}$ versus $1/C_{\mathrm{e}}$ indicates a good
correlation with the experimental data ($r=0.992$) and allows the
determination of the maximum biosorption capacity
$Q_{\max}=41.6$~mg${\cdot}$g$^{-1}$ as well as the Langmuir constant
$b=1.72$~L${\cdot}$mg$^{-1}$ (Figure~\ref{fig8}).  The equilibrium
factor $R_{\mathrm{L}}$ values were estimated from 0.037 to 0.104
depending on initial nitrate concentration. These  $R_{\mathrm{L}}$
values are significantly lower than 1 (close to zero), indicating that
the biosorption of nitrate by \textit{P.~oceanica} powder is highly
favorable (and tends toward irreversibility).

\begin{figure}
\includegraphics{fig07}
\caption{\label{fig7}Isotherm of nitrate biosorption onto
\textit{P.~oceanica} powder. (pH ${=}$\ 6, biomass dose ${=}$\ 0.167
g${\cdot}$L$^{-1}$, contact time ${=}$\ 60 min, temperature ${=}$\ 30
${\pm}$ 2~\textdegree C).}
\end{figure}

\begin{figure}
\includegraphics{fig08}
\caption{\label{fig8}Langmuir linear model of the isotherm of nitrate
biosorption onto \textit{P.~oceanica} powder. (pH ${=}$\ 6, biomass dose
${=}$\ 0.167 g${\cdot}$L$^{-1}$, contact time ${=}$\ 60 min, temperature
${=}$\ 30 ${\pm}$ 2~\textdegree C).}
\end{figure}

In the Freundlich model, the linear plot of Ln($Q_{\mathrm{e}}$) versus
Ln($C_{\mathrm{e}}$) allows for the estimation of the Freundlich
constant $K_{\mathrm{F}} = 26.98~\mathrm{mg}^{(1-(1/n))}{\cdot}
\mathrm{L}^{1/n}{\cdot}\mathrm{g}^{-1}$ and the value of $n=5.26$
(Figure~\ref{fig9}). The Freundlich modeling also shows a good
correlation  with experimental data ($r=0.983$) with an $n$ value
between 2 and 10, indicating a good biosorption process of nitrate by
\textit{P.~oceanica} powder. 

\begin{figure}
\includegraphics{fig09}
\caption{\label{fig9}Freundlich linear model of the isotherm of nitrate
biosorption onto \textit{P.~oceanica} powder. (pH ${=}$\ 6, biomass dose
${=}$\ 0.167 g${\cdot}$L$^{-1}$, contact time ${=}$\ 60 min,  temperature
${=}$\ 30 ${\pm}$ 2~\textdegree C).}
\end{figure}

Figure~\ref{fig10} shows the D--R linear form model allowing the
estimation of the maximum biosorption capacity 
$Q_{\max}=75.0$~mg${\cdot}$g$^{-1}$ and the activity coefficient $\beta
= 2\cdot 10^{-9}$ mol\tsup{2}${\cdot}$J$^{-2}$. The D--R model also
shows a good compatibility with experimental data ($r=0.988$), and
through the $\upbeta$ value it allows the calculation of the
biosorption mean free energy  $E=15.81$ kJ${\cdot}$mol$^{-1}$.

\begin{figure}
\includegraphics{fig10}
\caption{\label{fig10}D--R linear model of the isotherm of nitrate
biosorption onto \textit{P.~oceanica} powder. (pH ${=}$\ 6, biomass
dose ${=}$\ 0.167 g${\cdot}$L$^{-1}$, contact time ${=}$\ 60 min,
temperature ${=}$\ 30 ${\pm}$ 2~\textdegree C).}
\end{figure}

The parameters deduced from the Langmuir, Freundlich, and D--R models
are summarized in{\unskip\break} Table~\ref{tab3}. 

%tab3
\begin{table*}
\caption{\label{tab3}Langmuir, Freundlich, and D--R parameters for the
biosorption of nitrate onto \textit{P.~oceanica} powder}
\begin{tabular}{ccccccccc}
\thead
\multicolumn{3}{c}{Langmuir} & \multicolumn{3}{c}{Freundlich} & 
\multicolumn{3}{c}{D--R} \\\cline{1-3}\cline{4-6}\cline{7-9}
\parbox[t]{3pc}{\centering $b$ (L${\cdot}$mg$^{-1}$)} & 
\parbox[t]{3pc}{\centering $Q_{\max}$ (mg${\cdot}$g$^{-1}$)} & $r$ & 
\parbox[t]{7.5pc}{\centering $K_{\mathrm{F}}$ ($\mathrm{mg}^{(1-(1/n))}{\cdot}\mathrm{L}^{1/n}{\cdot}\mathrm{g}^{-1}$)} 
& $n$ & $r$ & \parbox[t]{3pc}{\centering $Q_{\max}$ (mg${\cdot}$g$^{-1}$)} & 
\parbox[t]{3pc}{\centering $E$ (kJ${\cdot}$mol$^{-1}$)}\vspace*{2pt} & $r$ \\
\endthead
1.72 & 41.6 & 0.992 & 26.98 & 5.26 & 0.983 & 75.0 & 15.81 & 0.988
\botline
\end{tabular}
\tabnote{pH ${=}$\ 6, biomass dose ${=}$\ 0.167 g${\cdot}$L$^{-1}$, contact
time ${=}$\ 60 min, temperature ${=}$\ 30 ${\pm}$ 2~\textdegree C.}
\vspace*{-2pt}
\end{table*}

The Langmuir model has the highest correlation coefficient value ($r$)
and fits better the experimental data. Moreover, although all three
models demonstrate good fit with experimental data based on the
Chi-square test (Table~\ref{tab4}), the Langmuir model exhibits the
smallest deviations between observed and theoretical values (smallest
value of $\rmchi_{\mathrm{obs}}^{2}$). Thus, biosorption mainly occurs
by a monolayer process on a homogeneous surface with a finite number of
identical sites, and no interaction between sorbed nitrate ions.
Furthermore, the value of mean free energy, calculated by D--R model,
is between 8 and 16 kJ${\cdot}$mol$^{-1}$, substantiating that the
nitrate biosorption was conducted by a chemisorption mechanism. In
accordance with this mechanism, the nitrate ions were mainly retained
by an ion exchange process as suggested in Section~\ref{sssec311} of pH
effect. They are also held by ion sharing between aqueous solution and
biosorbent surface as mentioned by previous research works using
\textit{P.~oceanica} to remove ammonium ions~\cite{68} and
metal-complexed textile dye~\cite{73} from aqueous solutions.

%tab4
\begin{table}
\caption{\label{tab4}Conformity Chi-square test applied at 95\%
confidence level to the isotherm models used}
\tabcolsep4pt
\begin{tabular}{cccc}
\thead
Chi-square test & Langmuir & Freundlich & D--R \\
\endthead
$\rmchi_{\mathrm{obs}}^{2}$ & 0.216 & 0.369 & 6.485 \\

\parbox[t]{4pc}{\centering $\rmchi_{\mathrm{c}}^{2}$\Lbreak
$n-1=18$; $\upalpha=0.05$} & 28.869\0 & 28.869\0 & 28.869\0\vspace*{3pt} \\

\parbox[t]{5pc}{\centering Retained hypothesis} & $\mathrm{H}_{0}$ & 
$\mathrm{H}_{0}$ & $\mathrm{H}_{0}$\vspace*{2pt}
\botline
\end{tabular}
\end{table}

Figure~\ref{fig11} displays the Langmuir nonlinear model superposed on
the experimental isotherm of nitrate biosorption. It is clearly
demonstrated that the Langmuir saturation plateau begins at an
equilibrium concentration of 6 mg${\cdot}$L$^{-1}$ of 
$\mathrm{NO}_{3}^{-}-\mathrm{N}$, corresponding to the initial
concentration of 12.5 mg${\cdot}$L$^{-1}$ of
$\mathrm{NO}_{3}^{-}-\mathrm{N}$. This initial concentration, in turn,
matches $12.5\times 4.428=55.35$ mg${\cdot}$L$^{-1}$ of
$\mathrm{NO}_{3}^{-}$. What is worthy to note is that the Langmuir
slope is high for concentrations below 55.35 mg${\cdot}$L$^{-1}$ of
$\mathrm{NO}_{3}^{-}$,  and much smaller for those above this
threshold. This observation clearly shows that the first biosorption
layer is saturated for an initial nitrate concentration of
55.35~mg${\cdot}$L$^{-1}$ and that the additional biosorption of nitrates
to form a second layer is practically difficult or very unlikely. It
indicates that the \textit{P.~oceanica} powder is efficient to treat
moderate content nitrate waters, such as most natural waters, potable
water, and secondary treated urban wastewater.

\begin{figure}
\includegraphics{fig11}
\caption{\label{fig11}\fontsize{9.7}{12}\selectfont
Isotherm and Langmuir non linear model of nitrate
biosorption onto \textit{P.~oceanica} powder (pH ${=}$\ 6, biomass
dose ${=}$\ 0.167 g${\cdot}$L$^{-1}$, contact time ${=}$\ 60 min,
temperature ${=}$\ 30 ${\pm}$ 2~\textdegree C).} 
\end{figure}

Furthermore, based on Langmuir and experimental maximum biosorption
capacity (41.6 and 39~mg/g, respectively), the comparison with other
nitrate adsorbents in the literature reveals that \textit{P.~oceanica}
powder is one of the most efficacious nitrate biosorbent
(Table~\ref{tab5}). This biosorbent also presents more advantages such
as being free, widely abundant in Mediterranean coasts, and easily
applied without previous preparations. The high value of nitrate uptake
by \textit{P.~oceanica} powder was reached without the need of any
chemical modification or cross-linking with other material as mentioned
in many previous research works (see Table~\ref{tab5}). At full scale,
the raw \mbox{water} to be denitrified by \textit{P.~oceanica} powder does not
require special treatment before or after biosorption, which makes for
an easy application and low-cost process. Besides, the biosorbent
saturated after use can be incinerated and stocked in confined places
far from natural water resources.

\begin{table*}
\caption{\label{tab5}The maximum adsorption capacity of nitrate by
various adsorbents cited in literature}
\tabcolsep3pt\fontsize{9.5}{12}\selectfont
\begin{tabular}{cccc}
\thead
Adsorbent & 
\parbox[t]{11pc}{\centering $Q_{\max}$ (mg of nitrate/g) (original unit in the reference)} & 
\parbox[t]{6.5pc}{\centering $Q_{\max}$ (mg of nitrate nitrogen/g)}\vspace*{2pt} & 
Reference \\
\endthead
\morerows{1}{\parbox[t]{12pc}{\centering Zinc chloride (ZnCl\tsub{2})
treated coconut granular activated carbon}} &
Langmuir: 10.26 at 25~\textdegree C & 2.31
& \morerows{1}{\cite{74}} \\
& Langmuir: 10.37 at 10~\textdegree C &
2.34 & \vspace*{5pt}\\ 

\morerows{1}{\parbox[t]{14pc}{\centering Activated carbon from
pomegranate rind (treated at 300~\textdegree C)}} & Langmuir: 13.10 &
2.96 & \morerows{1}{\cite{29}} \\
& Langmuir: 11.61 & 2.62 & \vspace*{5pt}\\ 

\morerows{1}{Sugarcane bagasse vermicompost} & Langmuir: 15.77 & 3.56 &
\morerows{1}{\cite{26}} \\
& Sips: 17.19 & 3.88 & \vspace*{5pt}\\ 

\morerows{2}{\raisebox{-2pc}{Agriculture waste}} & &
\parbox[t]{6.5pc}{\centering Wheat straw charcoal: 1.1}\vspace*{2pt} & 
\morerows{2}{\raisebox{-2pc}{\cite{27}}} \\
& & \parbox[t]{6.5pc}{\centering Mustard straw charcoal: 1.30}\vspace*{2pt} & \\
& & \parbox[t]{6.5pc}{\centering Commercial~activated carbon: 1.22} & \vspace*{5pt}\\ 

\morerows{2}{Chitosan} & Experimental: 19 & 4.29 &
\morerows{2}{\cite{25}} \\
& Langmuir: 8.03 & 1.81 & \\
& Freundlich: 23.85 & 5.39 & \vspace*{5pt}\\ 

\morerows{4}{\parbox[t]{12pc}{\centering Various materials for nitrate adsorption}} 
& Activated carbon: 4.14 & 0.93 & \morerows{4}{\cite{30}} \\
& Sepiolite: 3.46 & 0.78 & \\
& Sepiolite activated by HCl: & & \\
& Experimental: 9.80 & 2.21 & \\
& Langmuir: 38.16 & 8.61 & \vspace*{5pt}\\ 

\parbox[t]{12pc}{\centering Zeolite-supported zero-valent iron nanoparticles} & Langmuir: 22.94 & 
5.18 & \cite{75} \vspace*{5pt}\\ 

\morerows{1}{Local clay} & Experimental: 5.10 & 1.15 &
\morerows{1}{\cite{76}}  \\
& Langmuir: 27.77 & 6.27 & \vspace*{5pt}\\ 

\morerows{1}{\parbox[t]{15pc}{\centering Polyethylene glycol/chitosan composites and polyvinyl
alcohol/chitosan composites}} & (PEG) Langmuir: 35.03 & 7.91 &
\morerows{1}{\cite{31}} \\
& (PVA) Langmuir: 50.68 & 11.44 & \vspace*{5pt}\\ 

\morerows{1}{\parbox[t]{12pc}{\centering Activated carbon coated by Fe\tsub{3}O\tsub{4} magnetic
nanoparticles}} & Langmuir: 57.1 at 20~\textdegree C & 12.89 &
\morerows{1}{\cite{32}}  \\
& Langmuir: 64.5 at 35~\textdegree C & 14.56 & \vspace*{5pt}\\ 

\parbox[t]{12pc}{\centering Magnetic Co\tsub{3}O\tsub{4}/Fe\tsub{3}O\tsub{4} doped polyaniline
nanocomposite} & Langmuir: 68.96 & 15.57 & \cite{77} \vspace*{5pt}\\ 

Chitosan hydrogel beads & Langmuir: 92.1 & 20.80 & \cite{78} \\ 

\morerows{1}{\parbox[t]{12pc}{\centering Magnetic amine-crosslinked biopolymer}} & Langmuir: 65.36
at 25~\textdegree C & 14.76 & \morerows{1}{\cite{33}} \\
& Langmuir: 102.04 at 45~\textdegree C & 23.04 & \vspace*{5pt}\\ 
\splittabular

\morerows{3}{Modified wheat residue} & Modified wheat residue:
& 29.12 & \morerows{3}{\cite{79}} \\
& Langmuir: 2.08 (mmol/g) & & \\      
& Raw wheat residue & & \\
& Langmuir: 0.02 (mmol/g) & 0.28 & \vspace*{5pt}\\ 

\morerows{1}{Magnetic zeolite/nanocomposite} & Experimental: 421.5 &
95.2 & \morerows{1}{\cite{80}} \\
& Langmuir: 500 & 112.9 & \vspace*{5pt}\\ 

\parbox[t]{12pc}{\centering Polymer modified with quaternary ammonium
groups} & Langmuir: 221.81 & 50.1 & \cite{81} \vspace*{5pt}\\ 

\morerows{3}{Carbon} & Powdered activated carbon: & 140.0 &
\morerows{3}{\cite{24}}\\
& Langmuir: 10 mmol/g & & \\ 
& Carbon nanotube: & 350.0 & \\   
& Langmuir: 25 mmol/g & & \vspace*{5pt}\\
                   
\morerows{1}{The present work} & & Experimental: 39 & \\
& & Langmuir: 41.6 & 
\botline
\end{tabular}
\vspace*{3pt}
\end{table*}

The nitrate removal from aqueous media by biosorption onto the dead
seagrass \textit{P.~oceanica} makes it possible to solve several
problems simultaneously in Tunisia. On the one hand, the nitrate
removal from drinking water protects public health against several
diseases, particularly among infants and vulnerable populations
(pregnant women, elderly, etc.). On the other hand, not only does it
help clearing coastal areas of massive brown piles of dead
\textit{P.~oceanica} but it also makes the landscape pleasant for
swimming and tourist activities, as well as facilitates the fishing
boats mooring and maneuvering in ports. In addition, the nitrate
removal in wastewater treatment plants makes it possible to avoid
discharging nitrogen-rich water into aquatic media (sea, rivers, or
lakes) and, consequently, avoids the phenomenon of eutrophication. The
latter risks may lead to an ecological imbalance due to the depletion
of dissolved oxygen. 

\subsection{Kinetic modeling and biosorption mechanism investigation}
\label{ssec33}

Table~\ref{tab6} displays the calculated coefficients
$k_{1}$ and $Q_{\mathrm{cal.1}}$ for the pseudo-first-order model, 
$k_{2}$ and $Q_{\mathrm{cal.2}}$ for the pseudo-second-order model, and 
$k_{\mathrm{p}}$ for the intraparticle diffusion model, along with the 
corresponding correlation coefficient ($r$) for each. For the
pseudo-first order model, except for only one value of $r$ (0.995), all
the others are relatively low, lying between 0.950 and 0.990.
Furthermore, the difference between the calculated and experimental
values of nitrate nitrogen amount adsorbed at equilibrium is very high.
The error between the experimental and theoretical data ranges from 41
to 85\%. As a result, the pseudo-first-order model fits the
experimental data poorly and cannot describe the biosorption.

%tab6
\begin{table*}
\caption{\label{tab6}Kinetic parameters for the biosorption of nitrate
onto \textit{P.~oceanica} powder\vspace*{-2pt}}
\begin{tabular}{lcccccc}
\thead
& & \multicolumn{5}{c}{Initial nitrate-nitrogen concentration
(mg${\cdot}$L$^{-1}$)} \\\cline{3-7} 
& & 5 & 7 & 10 & 12 & 14 \\ 
\endthead
\multicolumn{2}{l}{Experimental $Q_{\mathrm{e}}$ (mg${\cdot}$g$^{-1}$)} & 
24.5\0\0\0 & 29.0\0\0\0 & 34.8\0\0\0 & 37.7\0\0\0 & 42.1\0\0\0 \\ 

\morerows{2}{Pseudo-first-order model} & $k_{1}$ (min$^{-1}$) & 0.007
& 0.028 & 0.034 & 0.033 & 0.033 \\
& $Q_{\mathrm{cal.1}}$ (mg${\cdot}$g$^{-1}$) & 3.6\0\0 & 5.1\0\0 & 20.9\0\0\0 & 21.8\0\0\0 &
24.9\0\0\0 \\
& $r$ & 0.990 & 0.995 & 0.974 & 0.968 & 0.950 \\ 

\morerows{2}{Pseudo-second-order model} & $k_{2}$ 
(g${\cdot}$mg$^{-1}{\cdot}$min$^{-1}$) & 0.009 & 0.010 & 0.013 & 0.018 &
0.009 \\
& $Q_{\mathrm{cal.2}}$ (mg${\cdot}$g$^{-1}$) & 25.6\0\0\0 & 30.3\0\0\0 & 35.7\0\0\0 & 38.5\0\0\0
& 43.5\0\0\0 \\
& $r$ & 0.997 & 0.999 & 0.999 & 0.999 & 0.999 \\ 

\morerows{1}{Intraparticle diffusion} & $k_{\mathrm{p}}$
(mg${\cdot}$g$^{-1}{\cdot}$min$^{-0.5}$) & 0.210 & 0.690 &  0.796 & 0.554
& 1.011 \\
& $r$ & 0.987 & 0.991 & 0.997 & 0.981 & 0.998
\botline
\end{tabular}
\ttabnote{pH ${=}$\ 6, biomass dose ${=}$\ 0.167 g${\cdot}$L$^{-1}$,
temperature ${=}$\ 30 ${\pm}$ 2~\textdegree C.}
\vspace*{-3pt}
\end{table*}

For the pseudo-second-order model, the correlation coefficients for all
initial concentrations are extremely high, ranging between 0.997 and
0.999. Moreover, the difference between the calculated and experimental
values of nitrate-nitrogen amount adsorbed at equilibrium is minimal
(error between 2.1 to 4.5\%) (Figure~\ref{fig12}). These results reveal
a good agreement between the experimental and theoretical data, which
means that the biosorption is well described by the pseudo-second-order
model. Consequently, biosorption is mainly governed by a chemical
process involving ion exchange or anion sharing between nitrate and
biosorbent, as described by some previous research works using
\textit{P.~oceanica} for ammonium~\cite{68} and
orthophosphate~\cite{44} removal.

\begin{figure}
\includegraphics{fig12}
\caption{\label{fig12}The plots of the pseudo-second-order model of
nitrate biosorption onto \textit{P.~oceanica} powder. (pH ${=}$\ 6,
biomass dose ${=}$\ 0.167 g${\cdot}$L$^{-1}$, temperature ${=}$\ 30 ${\pm}$
2~\textdegree C).}
\end{figure}

On the other hand, the plot of intraparticle diffusion model displayed
in Figure~\ref{fig13} clearly shows the three stages of mass transfer
as mentioned in Section~\ref{ssec27}. For all initial nitrate-nitrogen
concentrations, the first stage ($t^{1/2}$ between 0 and 3.1) is
characterized by high slope and describes the film diffusion, in which
there is transport of nitrate from solution to the exterior surface of
the cell membrane of \textit{P.~oceanica}. The second stage of the
curve, with lower slopes ($t^{1/2}$ between 3.1 and 7.8), represents
the intraparticle diffusion, in which nitrate is transported from the
extracellular environment into the interior of the cell membrane of
\textit{P.~oceanica}. The third stage, with null slopes ($t^{1/2} >
7.8$) reflects the equilibrium state of nitrate biosorption. As a
result of this analysis, the presence of three distinguished stages in
the curves for all initial concentrations and the relatively high
correlation coefficients of the linearized intraparticle diffusion
section ($r$ between 0.981 and 0.998; see Table~\ref{tab6}) lead to
the conclusion that intraparticle diffusion model plays an important
role in the nitrate biosorption onto \textit{P.~oceanica}.

\begin{figure}
\includegraphics{fig13}
\caption{\label{fig13}The plots of intraparticle diffusion model of
nitrate biosorption onto \textit{P.~oceanica} powder. (pH ${=}$\ 6,
biomass dose ${=}$\ 0.167 g${\cdot}$L$^{-1}$, temperature ${=}$\ 30 ${\pm}$
2~\textdegree C).} 
\end{figure}

\subsection{Thermodynamic study}\label{ssec34}

This study gives insight into the description and estimation of the
behavior of the biosorption reaction, such as its spontaneity,
thermicity, and the variation introduced on the randomness state in the
biosorbent--solution system. The same experimental data used in
Section~\ref{sssec314} (temperature effect) were used to calculate
$\Delta {G}_{T}^{0}$ at each chosen temperature. The values of
$\Delta {H}_{T}^{0}$ and  $\Delta {S}_{T}^{0}$ were determined
through the slope and intercept of the plot
$\mathrm{Ln}(K_{\mathrm{C}})=f(1/T)$ (Figure~\ref{fig14}),
respectively. All thermodynamic results are summarized in\break
Table~\ref{tab7}. 

\begin{figure}
\includegraphics{fig14}
\caption{\label{fig14}The plot of $\mathrm{Ln}(K_{\mathrm{C}})$ versus
$1/T$. (Initial aqueous concentration  ${=}$\ 5~mg${\cdot}$L$^{-1}$, pH
${=}$\ 6, biomass dose ${=}$\ 0.167 g${\cdot}$L$^{-1}$, contact time ${=}$\ 60
min).}
\end{figure}

%tab7
\begin{table}
\caption{\label{tab7}Thermodynamic parameters for the biosorption of
nitrate onto \textit{P.~oceanica} powder}
\tabcolsep=2pt
\begin{tabular}{ccccc}
\thead
\parbox[t]{2pc}{\centering $T$ (\textdegree C)} & $K_{\mathrm{C}}$ & 
\parbox[t]{4pc}{\centering $\Delta {G}_{T}^{0}$ (kJ${\cdot}$mol$^{-1}$)} &
\parbox[t]{4pc}{\centering $\Delta {H}_{T}^{0}$ (kJ${\cdot}$mol$^{-1}$)} &
\parbox[t]{4.8pc}{\centering $\Delta {S}_{T}^{0}$ (J${\cdot}$mol$^{-1}{\cdot}$K$^{-1}$)}\vspace*{2pt} \\
\endthead
30 & 4.70 & ${-}$3.90 & \morerows{4}{16.15} & \morerows{4}{65.97} \\
35 & 5.02 & ${-}$4.13 & & \\
40 & 5.50 & ${-}$4.44 & & \\
45 & 6.27 & ${-}$4.85 & & \\
50 & 6.91 & ${-}$5.19 & &
\botline
\end{tabular}
\xxxtabnote{Initial aqueous concentration ${=}$\ 5~mg${\cdot}$L$^{-1}$, pH
${=}$\ 6, biomass dose ${=}$\ 0.167  g${\cdot}$L$^{-1}$, contact time ${=}$\ 60
min.}
\end{table}

As shown by Table~\ref{tab7} and Figure~\ref{fig14}, the negative sign
of $\Delta G_{T}^{0}$ at all chosen temperatures means that
biosorption is spontaneous, whereby no energy input from outside the
system is required. The positive value of $\Delta {H}_{T}^{0}$
indicates that the biosorption is endothermic in nature as it was
noticed previously (Section~\ref{sssec314}). Furthermore, the low value
of biosorption heat ($\Delta {H}_{T}^{0}<40$ kJ${\cdot}$mol$^{-1}$) 
obtained in this work indicates that nitrate was mainly retained by
electrostatic interactions and no covalent bond should take place, as
described and suggested by a previous \mbox{research} work using chitosan
hydrogel beads to remove nitrate from water~\cite{78}. The positive
value of entropy, in turn, means an increase in the randomness state of
the solid--liquid interface after nitrate retention, which indicates
some changes occurring in the internal structure of the
biosorbent~\cite{82}.

\subsection{FTIR characterization}\label{ssec35}

The aim of infrared characterization is to compare the spectra of
\textit{P.~oceanica} powder before and after incubation with nitrates
to locate the appearance of new bands or the modification in intensity
or shape of certain already-existing bands. To this end, attributing
bands to functional groups and chemical bonds requires the use of
reference work and the infrared correlation chart. Figure~\ref{fig15}
shows two superposed spectra of \textit{P.~oceanica} before and after
nitrate retention. 

\begin{figure}
\vspace*{3pt}
\includegraphics{fig15}
\vspace*{3pt}
\caption{\label{fig15}Infrared spectra of \textit{P.~oceanica} powder.
(a) Native \textit{P.~oceanica}; (b) loaded
\textit{P.~oceanica}--$\mathrm{NO}_{3}^{-}$.}
\vspace*{3pt}
\end{figure}

Since \textit{P.~oceanica} is a plant material consisting mainly of
cellulose and other less important constituents, such as proteins and
lignin, their characteristic bands are clearly visible on both spectra.
Indeed, the band around 3340~cm$^{-1}$ can correspond either to the
valence vibration of bound --O--H, or to the valence vibrations of --NH
in amines emanating from the plant's vegetal proteins~\cite{83,84}. The
band around 2920~cm$^{-1}$ corresponds to the --C--H alkane valence
vibration~\cite{85,86}, while the band around 1638~cm$^{-1}$ is
attributed to the N--H amine deformation vibration~\cite{83}, as well
as the H--O--H valence vibration of the absorbed water~\cite{85,86}.
The band around 1508~cm$^{-1}$ could correspond to the N--H valence
vibration in the amino acids of plant proteins $\mathrm{RNH}_{3}^{+}$,
$\mathrm{R}_{2}\mathrm{NH}_{2}^{+}$, and
$\mathrm{R}_{3}\mathrm{NH}^{+}$ as well as the deformation vibration of
N--H in primary amines~\cite{83}. The bands around 1272 and 1033
cm$^{-1}$ can correspond to the C--O valence vibrations of aryl ether
and alkyl ether, respectively. On the spectrum of the loaded
\textit{P.~oceanica}--$\mathrm{NO}_{3}^{-}$, the appearance of new
bands provides evidence for the insertion of nitrate ions into the
structure of plant fibers. The two bands around 1609 and 1366 cm$^{-1}$
are characteristic of the nitrate ion as shown in a previous
study~\cite{83}. Indeed, the band around 1609 cm$^{-1}$, which forms a
shoulder with the band at 1638 cm$^{-1}$ by changing its shape, can
correspond to the asymmetric valence vibration of the nitro
--NO\tsub{2} group. The band around 1366 cm$^{-1}$ may correspond to
the symmetrical valence vibration of the nitro --NO\tsub{2} group, as
illustrated by Scheme~\ref{sch1} (see Section~\ref{ssec29} in the Material
and methods section).\looseness=1

The small new band that appears on the
\textit{P.~oceanica}--$\mathrm{NO}_{3}^{-}$ spectrum around 824
cm$^{-1}$ may correspond to the out-of-plane N--H deformation vibration
of amines~\cite{83}. The appearance of this band could be explained by
interactions between nitrates and amines. Given that the nitrate ion
can be neither a Lewis acid nor a Lewis base, the hypothesis of its
retention through coordination bonds is unlikely to be accepted. It
then appears that nitrate ions are retained mainly via electrostatic
bonds with charged amines $\mathrm{RNH}_{3}^{+}$,
$\mathrm{R}_{2}\mathrm{NH}_{2}^{+}$, and
$\mathrm{R}_{3}\mathrm{NH}^{+}$. The hypothesis of intraparticle
retention by the diffusion of nitrates in the cavities of the cellulose
crystal lattice can also be considered~\cite{87}.

\begin{figure*}
\vspace*{-2pt}
\includegraphics{fig16}
\vspace*{-3pt}
\caption{\label{fig16}SEM photograph of \textit{P.~oceanica} powder
before (a) and after nitrate biosorption (b).}
\vspace*{-3pt}
\end{figure*}

\subsection{SEM and EDS characterization}\label{ssec36}

Figure~\ref{fig16} displays the SEM micrographs of native
\textit{P.~oceanica} and \textit{P.~oceanica}--$\mathrm{NO}_{3}^{-}$
obtained at 50 and  5~$\upmu$m magnifications. The micrograph at
50~$\upmu$m obviously shows the ultimate cellulosic fibers. At this
resolution level, while the micrograph of native \textit{P.~oceanica}
illustrates the condensed and assembled fibers, that of
\textit{P.~oceanica}--$\mathrm{NO}_{3}^{-}$ elucidates isolated fibers
that are far from each other. These observations can be explained by
the insertion of \mbox{nitrate} ions between the fibers and the formation of
new electrostatic bonds, including hydrogen bonds and van der Waals
forces. These new forces impose a sort of arrangement and orientation
of the fibers in well-defined  directions~\cite{88}, and thus the
fibers move away from each other. The micrograph at 5~$\upmu$m shows
the \mbox{surface} of one ultimate cellulosic fiber. At this resolution level,
the micrograph of native \textit{P.~oceanica} indicates a homogenous
surface covered by intact plant matter. However, the micrograph of
\textit{P.~oceanica}--$\mathrm{NO}_{3}^{-}$ suggests a degraded and
fragmented layer, which supports the hypothesis of chemical structure
modification on the fiber surface following interaction with nitrate
ions. 

Figure~\ref{fig17} shows the EDS spectra (chemical analysis) of native
\textit{P.~oceanica} and loaded
\textit{P.~oceanica}--$\mathrm{NO}_{3}^{-}$. The spectrum (a) of native
\textit{P.~oceanica} shows the two significant signals of oxygen and
nitrogen, which provides strong evidence not only that nitrogen exists
in the chemical structure of the plant biomass as amine groups, but
also that oxygen exists as alcohol groups and in the pyranosic rings of
cellulosic chains. The spectrum (b) of
\textit{P.~oceanica}--$\mathrm{NO}_{3}^{-}$ shows a notable increase in
nitrogen and oxygen signals, explained by the insertion of nitrate into
the plant's biomass structure. Moreover, it is worth noting that the
oxygen signal increases significantly more than that of nitrogen. This
observation can be explained by the fact that each incoming nitrate ion
introduces one nitrogen atom and three oxygen\break atoms.

\begin{figure*}
\includegraphics{fig17}
\caption{\label{fig17}EDS spectra of \textit{P.~oceanica} powder before
(a) and after nitrate biosorption (b).}
\end{figure*}

\subsection{Biosorption mechanism} 

Understanding the biosorption mechanism is crucial to identifying the
chemical groups or ions on the biosorbent surface that are involved in
nitrate retention. Kinetic study shows that the biosorption is well
described by the pseudo-second-order model rather than the
pseudo-first-order one, and the D--R isotherm modeling, in turn,
demonstrates that the mean free energy is between 8 and 16
kJ${\cdot}$mol$^{-1}$. Together, both results confirm that nitrate was
retained by chemisorption~\cite{89,90,91,92}. In accordance with this,
the nitrate ions were mainly retained by an ion exchange process, as
observed and described by some previous studies~\cite{68,73}. The
hypothesis of chemisorption via ion exchange with hydroxide ions was
experimentally proven during the pH effect study. In fact, following
the biosorption process, the final pH of the aqueous solutions was
consistently higher than the initial pH across the range of 5--12.
Furthermore, the intraparticle isotherm model reveals good agreement
with the experimental data, and its plot clearly displays the three
stages upon which its assumptions are based (see Figure~\ref{fig13}).
This behavior suggests that nitrate \mbox{retention} by \mbox{intraparticle}
diffusion may occur, but to a lesser extent than by ion exchange. On
the other hand, the analysis of infrared spectra indicates that nitrate
may interact through electrostatic forces with the charged amine on the
biomass surface ($\mathrm{RNH}_{3}^{+}$;
$\mathrm{R}_{2}\mathrm{NH}_{2}^{+}$; $\mathrm{R}_{3}\mathrm{NH}^{+}$).
However, at pH ${=}$\ 6 (optimal value for biosorption), the number of
protonated amines on the biosorbent surface is very small, which
suggests that this process of nitrate retention is very weak compared
to that of ionic exchange. In conclusion, it is estimated that at pH
${=}$\ 6, nitrate retention occurs predominantly through ion exchange,
which appears to be the most significant mechanism, while electrostatic
interactions and intraparticle diffusion may also occur but play a
considerably lesser role.

\section{Conclusion}\label{sec4}

The removal of nitrate from aqueous solutions was performed by
biosorption onto abundant sea plant \textit{P.~oceanica}. The
biosorption was carried out in batch experiments at room temperature
and under optimal conditions. It was found that the Langmuir model fit
better the experimental data, confirming a monolayer biosorption
process onto homogeneous surface and with a finite number of identical
active sites. The Langmuir maximum nitrate uptake was found to be equal
to 41.6~mg${\cdot}$g$^{-1}$ of nitrate nitrogen, showing good efficiency
of the biosorbent used. This high value of nitrate uptake by
\textit{P.~oceanica} powder was reached without the need for any
chemical modification or cross-linking with other material as mentioned
in many previous research works reported in the literature. The
Dubinin--Radushkevich modeling, in turn, gives a mean free energy of
15.81 kJ${\cdot}$mol$^{-1}$, illustrating that nitrate was retained
through a chemisorption process involving ion exchange or ion sharing
between the aqueous solution and the adsorbent surface. The
thermodynamic study shows that biosorption is spontaneous, endothermic
in nature, and accompanied by an increase in entropy. The spectroscopic
characterization by FTIR indicates that nitrate was mainly retained
through electrostatic interactions with protonated amines onto the
biosorbent surface. This result is further supported by EDS and SEM
analyses. Consequently, the sea plant \textit{P.~oceanica} demonstrates
strong potential as an effective biosorbent for nitrate removal from
natural waters, offering a sustainable alternative to conventional
denitrification methods. 

\section*{CRediT authorship contribution statement}

\textbf{Abdelbasset Ben Dkhil:} Conceptualization, Methodology, Investigation, Formal Analysis, Writing – original draft.

\noindent \textbf{Haitham Elleuch:} Investigation, Vizualisation.

\noindent \textbf{Anissa Somrani:} Writing – Review and Editing, Formal Analysis.

\noindent \textbf{Amor Hafiane:} Writing – Review and Editing, Project administration.

\section*{Declaration of interests} 
          
The authors do not work for, advise, own shares in, or receive funds
from any organization that could benefit from this article, and have
declared no affiliations other than their research organizations.

\section*{Funding}

The authors would like to thank the Ministry of Higher Education and
Scientific Research of Tunisia for financial support of this research
work.

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