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Article

Sustainable Synthesis of Hydro Magnesite Fire Retardants Using Seawater: Characterization, Yield Modeling and Process Optimization

by
Mohammad Ghaddaffi Mohd Noh
1,2,*,
Nor Yuliana Yuhana
2 and
Mohammad Hafizuddin bin Hj Jumali
3
1
PETRONAS Research Sd. Bhd, Lot 3288 & 3289 Off Jalan Ayer Itam Kawasan Institusi Bangi, Kajang 43000, Selangor, Malaysia
2
Department of Chemical and Process Engineering, Faculty of Engineering & Built Environment, Universiti Kebangsaan Malaysia, Bangi 43600, Selangor, Malaysia
3
Department of Applied Physics, Universiti Kebangsaan Malaysia, Bangi 43600, Selangor, Malaysia
*
Author to whom correspondence should be addressed.
Fire 2025, 8(10), 409; https://doi.org/10.3390/fire8100409
Submission received: 9 July 2025 / Revised: 22 August 2025 / Accepted: 23 September 2025 / Published: 21 October 2025
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Abstract

The Global Cement and Concrete Association (GCCA) estimated that by 2050, 36% industry-wide sustainable value will be created, which includes sequestering CO2 into the cement and concrete industry to produce commercially feasible high-value products. Direct utilization of CO2 in the cement and concrete industry, which utilizes natural and sustainable materials, is gaining momentum. Naturally occurring mixtures of hydro magnesite and huntite are important industrial minerals which, upon endothermic decomposition over a specific temperature range, will release water and CO2. This unique chemistry has led to such mixtures being successfully utilized as fire retardants, replacing aluminum hydroxide or Alumina Tri-Hydrate (ATH). Despite the developed marketplace for magnesium-based fire-retardant products, there is little mention of CO2 mineral carbonation methods, which attempt to recover and convert magnesium from natural seawater or industrial waste into oxides or carbonates as part of the carbon sequestration initiative. The hypothesis to be proven in this work states that if the process of seawater mineral carbonation is prematurely quenched, Mg2+ ionic species in seawater adsorbed on the calcite lattice formation will be trapped and therefore recovered in various oxidized forms, such as magnesium oxides, magnesium hydro magnesite, and magnesium carbonate precipitates. A novel method to recover magnesium Mg2+ ions from seawater was successfully explored and documented; as such, from an initial concentration of 1250 ppm Mg2+ in raw seawater, the average concentration of spent Mg2+ ions after the reaction was as low as 20 ppm. A very efficient near-total recovery of Mg2+ from the seawater into the solid precipitates was recorded. Subsequently, the process for continuous seawater mineral carbonation for the production of magnesium/brucite/huntite products was successfully proven and optimized to operate with a 30 s reaction time, a dynamic feedstock concentration, [CaO] at 1 gpl in seawater and a room temperature reaction temperature (30 °C), where the average yield of the fire-retardant magnesium-based compounds was 26% of the synthesized precipitates. Approximately 5000 g of the hydro magnesite materials was molded into a fire-retardant brick or concrete wall, which was subjected to an accredited fire performance and durability testing procedure BS476-22:1987. There were encouraging results from the fire resistance testing, where the fire-retardant material passed BS476-22:1987, with performance criteria such as physical integrity failure, the maximum allowable face temperature, and a minimum duration before failure, which was up to 104 min, evaluated.

1. Introduction

The International Energy Agency (IEA) [1] has stated that, in order for the energy sector to achieve net zero emissions by 2050, the global scale of Carbon capture and Sequestration, CCS, in the years 2030 and 2050 must, respectively, be 10–15 times and 100 times greater than the current 40 Mtpa as of 2020. In the wake of such a global uprise in the carbon dioxide CO2 emissions level, much effort has been put into maturing CCS [2] and Carbon Capture, Utilization and Sequestration, CCUS [3,4,5], technology portfolios. CCS focuses more on upscaling the science and technology of capturing CO2, injecting and trapping it back into the Earth’s geological storage. CCUS, on the other hand, has an added dimension of CO2 utilization, such as reusing it directly as a green solvent, industrial coolant, or even as dried ice in the food and beverage industry, or indirectly by converting it into other forms of molecules with higher value or usage. Enobong Hanson et al. [5] provided a comprehensive examination of CCUS technologies, focusing on their advancements and future prospects, and highlighting recent technological improvements and associated challenges. According to Sakakura et al. [6], there are four major methodologies for transforming CO2 into useful chemicals, such as the usage of high-energy starting materials such as hydrogen, unsaturated compounds, small-membered ring compounds, and organometallics. However, SF Chang et al. [7] commented that techno-economic and environmental evaluations for CCUS have only been conducted in a limited number of relevant studies, which tend to evaluate individual capture technologies across various industries, including the building and construction material industry. Moving forward, the Global Cement and Concrete Association (GCCA) [8] estimated that by 2050, its CCUS mitigation proposals have the potential for 36% industry-wide value creation, which includes CO2 reduction initiatives such as sequestering CO2 into the cement and concrete industry to produce commercially feasible high-value products. As such, the development of a large-scale carbonation process technology using sustainable materials such as seawater or industrial waste to produce a wide range of high-value building materials is hereby proposed. The sequestration of CO2 into these sustainable building and construction materials will potentially contribute to some 6% of the global carbon utilization in the journey to achieve net zero emissions by 2050.

2. State of the Art

LeClerc et al. (2025) [9] reported technological breakthroughs that largely began as early as the 1990s, which have now progressed into a diverse and plentiful portfolio of technological and scientific advancements and commercialized technologies and product portfolios related to CO2 utilization, as shown in Figure 1a,b below.
As pointed out by Aresta et al. [10], CO2 is considered a “spent” molecule of carbon C whereby only nature is able to naturally recycle it into thousands of naturally occurring useful carbon-based compounds via biological conversion pathways available within the microbial and plant kingdom using the photosynthetic energy from the sun. In view of the energy standpoint [11], conversion of CO2 will require a certain amount of energy depending on the downward steps of the state of carbon oxidation from +4 in CO2 to that of the target product. Consequently, the reactions of CO2 can be divided into two main categories which are firstly, the low energy processes, in which C maintains its +4 oxidation state or is lowered by only one unit (a C-C bond is formed) which bear to the formation of carboxylates, carbamates, carbonates, urea, polymeric materials such as polycarbonates and polyurethanes, inorganic carbonates, and also hydrogen carbonates. The second groupings of CO2 reactions involves high energy processes, in which C goes down to oxidation states lower by at least two units below +4: HCOOH, CO, H2CO, CH3OH, CH4, and hydrocarbons where the external input of energy is needed and can be delivered from several energy carriers such as electrons (electrochemical reduction), hydrogen (hydrogenation reactions), metals (reaction with elemental Group 1 metals), and also from plasma and radiations. Due to the limitation of available excess energy in large amounts to convert CO2 to either one of the high-energy molecules mentioned in the second groupings above, large-scale conversion of CO2 to fuels is still not economically scaled up globally.
One interesting option is to look more closely into the large-scale conversion of CO2 to carbonates, which is not relatively energy-intensive compared to end products like long carbon chain fuels and energetic chemicals. While successful conversion of CO2 to calcium carbonate products has been demonstrated by many, aragonite synthesis and upscaling, which is another crystalline morphology of calcite, is still very limited and requires large amounts of chemical inducers such as magnesium chloride. As mentioned earlier, the Global Cement and Concrete Association (GCCA) [8] estimated that by 2050, its proposed CO2 reduction initiatives will mostly be applied to the cement and concrete industry. Once captured, CO2 can be used commercially as green feedstock in high-value products. According to Sarah Danelli et al. [12], Portland Cement (PC), being the most used manufactured material globally, has an estimated 4.1 billion tons of annual production volume will be the main target in mitigating the high CO2 emissions in the cement industry. Carbon capture in the construction industry can be demonstrated via the mineral carbonation of PC derivative products, such as in the form of integral structures, post-construction, or demolition waste (CDW), which is usually referred to as CO2 sinks. Most of the commercialized CO2 mineralization in the cement and concrete industry produced improved materials for use, such as supplementary cementitious materials SCM or filler materials, magnesium carbonate enriched with silica for use in concrete, synthetic aggregate and ACT, synthetic aggregates end products, raw material for clinker SCM, lower carbon footprint ready mix concrete, and precast materials [12].
Direct utilization of CO2 in the cement and concrete industry utilizing natural and sustainable materials is also gaining momentum. In this research and commercial space, Carbon Upcycling Technologies (CUT) [13,14,15,16] capitalized the patented Mechanically Assisted Chemical Exfoliation (MACE) process to transform alternative feedstocks including fly ash, silicates, aggregate fines, steel slags, glass, biomass, volcanic rocks, natural pozzolans, metal mine tailings, and clays, among others, into enhanced materials by combining CO2, either low or high purity, with industrial by-products or natural materials in a catalytic reaction system. Paebbl et al. also developed a process that involves the comminution of silicate rocks, mainly olivine, which accelerates the absorption of CO2 in a batch process into silica-enriched magnesium carbonate [17,18]. In another development, Carbon8 Systems in 2006 commercialized their patented process of Accelerated Carbonation Technology (ACT) [19,20], which solidifies industrial residues such as ash and cementitious waste into a hardened pellet form that can be used as a substitute for natural aggregate in concrete. Startups such as Blue Planet Systems [21] commercialized the process known as Geomimetic [22,23,24,25], which intends to simulate the natural mineralization of carbonate rocks or limestone by accelerating on an industrial scale, while Neustark, founded in 2019, captures and mineralizes CO2 into CDW through an accelerated process of mineralization, where the CO2 turns into limestone industrial products [26].
Another area of direct utilization of CO2 is producing cement precast with carbon-activated residue. Carbstone Innovation, in 2004, developed a process to accelerate the carbonation of slag residue to produce precast materials such as tiles, roof tiles, paving bricks, kerb stones, and building block products [27,28,29], where it is estimated that 1 m3 of Carbstone bricks stores 350 kg of CO2 [30]. Carbicrete, founded in 2016, commercialized a cement-free binder carbon-activated steel slag [31,32], where each Carbicrete masonry unit is estimated to avoid 3 kg of CO2 [33]. CarbonBuilt has a patented a concrete mix design and its curing process where in this novel formulation, some 60 to 80% of portlandite cement is substituted by Ca(OH)2, low-carbon SCM, and fillers [34], while in another development, Solidia Technologies patented a technology [35] which modifies the cement formulation by creating a kind of carbonate calcium silicate cement (CCSC) where reduced CO2 emissions from calcium carbonate breakdown and CO2 absorption during curing leads to a reduction in the carbon footprint from 810 kgCO2/ton of Portland cement clinker to 565 kgCO2/ton of calcium silicate-based clinker [36,37]. Fortera launched ReCarb in 2019, which intercepts CO2 emitted during limestone calcination and converts it into low-carbon cement known as ReAct (Reactive Calcium Carbonate) during cement hydration, increasing the strength of the matrix [38]. CO2 injection in ready-mix concrete is also a form of direct CO2 utilization. CarbonCure, in 2012, injected previously captured CO2 into fresh concrete during the mixing to be used as ready mix or precast, where CO2 combines with the calcium ions available in the concrete mix, to save between 7 and 11 kg of CO2 per cubic meter of ready-mix concrete [39].
In the specialty product space, there are many producers utilizing magnesium hydroxide or brucite for flame-retardant applications. Naturally occurring mixtures of hydro magnesite and huntite are important industrial minerals which, upon endothermic decomposition over a specific temperature range, will release water and CO2, consequently leading to such mixtures being successfully utilized as fire retardants in replacement of aluminum hydroxide or Alumina Tri-Hydrate (ATH). It has been reported by many that mixtures of natural hydro magnesite and huntite have been commercially exploited as fire retardants since the late 1980s. Toure [40] has published his work detailing the usage of mixtures of hydro magnesite and huntite as a fire retardant in an ethylene propylene copolymer. Several studies conducted by [41,42,43,44] have illustrated the effectiveness of magnesium hydroxide as a flame retardant in plastic and concluded that magnesium hydroxide is effective at reducing smoke emissions from burning plastics. Hollingbery et al. [45] studied the decomposition mechanism and morphology of mixtures of hydro magnesite and huntite and found several advantages as hydro magnesite starts to decompose at a higher temperature than aluminum hydroxide, allowing usage in polymers with higher molding temperatures [46]. Also, the wider endothermic decomposition range of hydro magnesite and huntite provides a cooling effect over a wider temperature range than aluminum hydroxide or magnesium hydroxide alone. Additionally, the platy morphology of the huntite will provide a barrier to the transport of combustible decomposition products to the flame while also reinforcing the char layer and providing further protection to the underlying polymer by endothermic decomposition in response to high external heat fluxes. Huntite on its own is beneficial to polymers like Polyether Ether Ketone PEEK where the polymer’s high melt and decomposition temperatures mean that aluminum hydroxide or magnesium hydroxide would be unsuitable. In the same study, authors L. A. Hollingberya and T. R. Hull (2010) [47] deduced that hydro magnesite has a similar decomposition enthalpy to aluminum hydroxide, although the heat capacity of aluminum hydroxide is slightly higher than hydro magnesite itself.

3. Theory

Despite the developed marketplace for magnesium-based fire-retardant products, there is little mention of CO2 mineral carbonation methods, which extend its feedstock coverage to recover and convert magnesium from natural seawater or industrial waste into oxides or carbonates as part of the carbon sequestration initiative. This is due to the fact that the carbonate of magnesium is relatively more soluble than calcium carbonate, posing a drawback in maximizing solid product recovery.
Classical mineral carbonation using seawater, as reported by Dami Kim et al. (2016) and Hsing-JH et al. (2023) [48,49], suggests that the equilibrium governing Equations (1)–(9) is proposed for the reaction steps for the CO2 carbonation process for aragonite synthesis using magnesium as a chemical inducer is as mentioned below:
Ca(OH)2(s) + MgCl2 (aq)[from seawater] → Mg(OH)2(s) + CaCl2(aq)
Ca(Cl)2(aq) + H2CO3(aq) + Mg(OH)2(s) → CaCO3(s) + MgCl2(aq) + 2H2O
CO2 absorption
CO2(g) → CO2(aq)
Hydration of dissolved CO2(aq)
CO2(aq) + H2O ←→ H2CO3(aq)
Reaction of hydroxyl ions with dissolved CO2 in the case of high pH of the solution
CO2(aq) + OH−(aq) ←→ HCO3−(aq)
HCO3−(aq) + OH−(aq) → CO32−(aq) + H2O
Dissociation of calcium hydroxide
Ca(OH)2(aq) → Ca2+ + 2OH−(aq)
Ca(OH)2(s) + Mg2+ + 2Cl− (aq) → Mg(OH)2(s) + Ca2+(aq) + 2Cl−(aq)
Complete Precipitation
Ca(Cl)2(aq) + CO32−(aq) + H2O+ Mg(OH)2(s)→ CaCO3(s)+ MgCl2 (aq)+ 2H2O
In terms of reaction kinetics, carbonate precipitation reactions are characterized by a timescale as proposed by Mitchell et al. [50]. Under well-mixed conditions (i.e., free of mass transport limitations) at 25 °C and 1 atm, the equilibrium described by reaction 8 occurs within t = 5.0 × 10−11 s.
CO2(g) <==>CO2(aq)
The aqueous species H2, CO3, HCO3−, and CO32−, as described by reactions (5)–(7), reach equilibrium within 10−2 s.
CO2(aq) + H2O <==> H2CO3 + HCO3−
H2CO3 <==> H+ + HCO3−
HCO3− <==> H+ + CO32−
However, the equilibrium with respect to Ca2+ (i.e., reactions (8) and (9)) is only reached in 103 s, which is roughly between 10 and 20 min.
Ca2+ + CO32− <==> CaCO3(aq)
CaCO3(aq) <==> CaCO3(s)
Firstly, in building up the hypothesis of this work, we consider the study by Mitchell et al. [50], which proposed a kinetics model where calcium carbonate precipitation is susceptible to relatively slow formation kinetics. Secondly, from the molecular simulation work of WenHao S. et al. [51], the effect of Mg2+ incorporation on the calcite surface energy was found to drastically inhibit calcite nucleation. Surface energies are obtained from Discrete Fourier Transform DFT computations of hydrated calcium carbonate surface slabs with partial Mg2+ substitution on calcium sites, as shown in Figure 2. The kinetic phase diagram study showed that for the ratio of Mg/Ca = 5.2 (modern seawater), only aragonite is preferred to nucleate, while concurrent nucleation of calcite and aragonite occurs for a broad span of supersaturations near Mg/Ca = 2, which potentially yields brucite or hydro magnesite.
Thirdly, this brings us to the most commonly expressed idea is that magnesium Mg2+ inhibits the growth of calcite, therefore allowing aragonite phase to crystallize rapidly as pointed out by Bischoff, 1968 [52]; Bischoff and Fyfe, 1968 [53]; Berner, 1975 [54]; Fernández-Díaz et al., 1996 [55]; De Choudens-Sánchez and González [56]. Inhibition of calcite growth is generally attributed to “poisoning” of the growth sites on the calcite crystal faces by Mg ions that are more strongly hydrated than the Ca ions [57]. Henceforth, our hypothesis in this work states that if the process of seawater mineral carbonation is prematurely quenched, Mg2+ ionic species which adsorbed on the calcite lattice formation will be trapped and therefore recovered in various oxidized forms such as magnesium oxides, magnesium hydro magnesite, and magnesium carbonate precipitates.
Hence, taking Equation (7) for the case of premature quenching of the seawater mineral carbonation reaction, or the incomplete precipitation of aragonite crystalline particles, Mg2+ adsorbed on the crystalline lattice of CaCO3(s) will be recovered to form a family of magnesium based solid precipitate such as hydro magnesite, huntite, and brucite as shown in Equations (14)–(16) below.
Incomplete Precipitation:
Brucite
Ca(Cl)2(aq) + CO32− (aq) + 2H2O + Mg(OH)2(s)(premature)→Mg(OH)2(s) + CaCO3(s) + MgCl2(aq)+ 2H2O
Hydro magnesite
(Ca(Cl)2(aq) + 5CO32−(aq) + 6H2O + 6Mg(OH)2(s)(premature)→ Mg5(CO3)4(OH)2·4H2O(s) + CaCO3(s) + MgCl2(aq)+ 2H2O
Huntite
2Ca(Cl)2(aq) + 5CO32−(aq) + H2O + 4Mg(OH)2(s)(premature) → Mg3Ca(CO3)4(s) + CaCO3(s)+ 2MgCl2(aq) + 2H2O

4. Method

4.1. Experimental Setup and Procedure Using Lab Scale Glass Reactor

This setup is to investigate the effect of solution temperature, ratio of lime CaO in seawater, and reaction time to maximize magnesium complex compounds, such as hydro magnesite, huntite, and brucite yield, with the expected results of establishing a correlation between said factors and hydro magnesite/huntite/brucite yield, and to select optimum operating conditions for a continuous reactor system.

4.2. Feedstock Preparation

Solution of seawater was prepared using a 33 g/L blend of red sea salt [58] dissolved into deionized water. At this salt concentration, the Ca2+ ion was in the range of ~400 ppm, and the Mg2+ ion in the range of ~1100 ppm, with solution alkalinity between pH 2.2 and pH 2.4. Then, CaO was blended into this seawater blend at 1 g/L, 3 g/L, and 5 g/L to form a slurry feedstock, where subsequently the feedstock was heated to various target temperatures (30–90 °C). CO2 feed gas with a constant flow rate of 300 mL/min was bubbled into the reactor at different reaction time ranges (1–60 min). The solution reacted and formed a whitish precipitate. After the experiment, the liquid and solid were separated via vacuum filter, and the solid sample was analyzed using equipment such as Hitachi FESEM SU8020 (Hitachi High-Tech Group, Japan. Origin of country Japan, sourced from Kajang, local agent Malaysia) for FESEM-EDX and SEM imaging, XRD Bruker D8 Advance (Bruker Physik-AG in Karlsruhe, Germany) for X-Ray Diffraction analysis, Induction Coupled Plasma ICP-OES Avio 500 (PerkinElmer, Waltham, MA, USA) for ionic content analysis, and TGA STARe System TGA/DSC 1 HT/1600 (Mettler-Toledo International Inc., Greifensee, Switzerland)analytical methods for thermal gravimetric analysis.

4.3. Experimental Setup and Procedure Using Continuous Reactor

The experimental work was carried out in the continuous mineral carbonation reactor shown in Figure 3 below. This setup intended to validate the performance of the optimum operating conditions using the continuous flow method of brucite production.
The procedure for starting up and operating the continuous reactor is provided. Pressure reactor vessel 800 mL was filled with compressed CO2 up to 5 bar. The feedstock solution consisted of lime CaO and seawater was prepared at room temperature, where the concentration of magnesium was approximately 1200 ppm, or using 33 g/L of red sea. The feedstock solution was sprayed into the reactor at 40 mL/min for 3 min before the manual valve was opened. The outlet valve and flow rate were maintained at 40 mL/min to ensure 3 min reaction time inside the reactor. The slurry product was collected at the bottom. Reacted products with whitish–soapy presence were sampled for further analysis. The experiment was repeated for differential conditions such as the reaction temperature at room temp RT of 30 °C and high temperature of 90 °C; CO2 contact time of 3 min and 60 min; and lime CaO concentration in seawater of 1 g/L, 3 g/L, and 5 g/L. Solid precipitate samples were dried and tested for the physical morphology via XRD, while the filtered liquid underwent ICP testing to establish the concentration of dissolved metal ions.

5. Design of Experiment

5.1. Design of Experiment

The Two-Level Full-Factorial Design was chosen as the design of the experiment methodology. The test consists of 24 runs: three parameters, two levels (high and low) with three replicates, totaling 24 runs. As shown in Figure 4 below, the parameters under investigation are as follows: (1) temperature: 30–90 °C, (2) concentration of CaO in seawater: 1 g/L and 5 g/L, and (3) reaction time: 3 and 60 min. The provisional equation for hydro magnesite, brucite, and huntite yield is as follows:
Yield (hydro magnesite, brucite, and huntite) =
α_1 × A + α_2 × B + α_3 × C + α_4 × AB + α_5 × BC + α_c × AC + α_7 × ABC

5.2. Results

The results of magnesium/brucite/huntite yield are tabulated in Table 1 below for the respective parameters under investigation, where the response of aragonite percentage yield recorded was between 0% and 85.0%
The standard run and the response are visualized in a 3D factor cube to better picture the responses: individual single-factor responses, two-way factor interaction responses, and three-way factor interaction responses. A first-cut understanding of how the three factors interact, whether individually or interactively, is summarized below:
Point_1 provides the baseline response when all factors are reset to low, in which if all the factors are lowered, the magnesium complex yield Y has a sum of response of 70.0 + 70.0 + 70.0 = 210.0. If only factor 1, which is temperature, is increased to 90 °C from baseline 30 °C, and point 2 is of interest, we see that the magnesium complex yield Y decreases to a sum of 14.9 + 14.9 + 14.9 = 44.7, or in terms of percentage response, by increasing temperature by a quantum of 60 °C, the aragonite yield decreases by 78.7%. Hence, temperature has a major negative correlation effect. The effect of reaction time can be seen at Point_3 with sum of response of 2 +2 + 1 = 5. Hence, the increase in reaction time by another 9 min drastically decreases or nullifies the magnesium complex yield by 97.6%. At Point_5, the sum effect of increasing the concentration of reactant CaO by 5-fold from 1 gpl to 5 gpl marginally increases the magnesium complex yield to only 251 from 210, or only increases by 20%, which is a positive effect. Hence, the effect of increasing the temperature and reaction time should be avoided, as it provides a significantly negative correlation to the yield of magnesium complexes. The aragonite yield associated with the factors under investigation is provided in Figure 4 below and Table 2 below.

5.3. Investigation of Main Factors for Magnesium Complex Compounds Yield

The individual and compounded effects of the factors investigated: Temperature A, Duration B, and CaO Concentration C. The experimental results are tabulated and analyzed using Design Expert to determine the main effects associated with the factors A, B, and C and the interactions between them AB, BC, and ABC. The half-normal plot, as shown in Figure 5 below, suggests that the effect of B—duration is the most significant, while A—temperature and AB are the next significant factors. C—concentration of CaO is not a significant factor in terms of increasing the magnesium complex compound yield.
Hence, experimental duration, reaction temperature, and the two-way factor interaction of AB (temperature x duration) have a significant effect on the yield of magnesium complex compounds. Both duration and temperature have a negative effect, while the two-way interaction factor AB has a positive effect. Hence, the general equation of the yield is adjusted or simplified to reflect only the components that significantly affect the magnesium complex compounds yield, as such:
Aggregate hydro magnesite, brucite, and huntite yield Y = α_1 A+α_2 B +α_4 AB

5.4. Optimization of Aggregate Hydro Magnesite, Brucite, and Huntite Formation in Seawater Media

The optimization of the aggregate hydro magnesite, brucite, and huntite yield was conducted along the values of the reaction time factor at 1 min, 5 min, 10 min, and 30 min, with the temperature factor kept at a low 30 °C. When the reaction time was reduced from 3 min to 1 min, the yield increased from 14.9% to 25.7% at a CaO concentration of 1 gpl. The values of temperature were kept at 30 °C and the concentration of CaO at 1 g/L during this optimization at other points associated with reaction times of 5 min, 10 min, and 30 min. The optimization planning is graphically represented in Figure 6 below.
Apart from optimizing the effect of reaction time, as in Figure 6, extended experimentation to validate the effect of CaO concentration was conducted to include the following ratio of Mg/Ca associated with feedstock blends of CaO in seawater of 1 g/L, 3 g/L, and 5 g/L to investigate the kinetics of magnesium complexes formation under different Mg/Ca ratios or concentrations.
With reference to the trend shown in Figure 7a–c, the yield of magnesium-based compounds reduced drastically in a logarithmic down-trending after the initial 5 min reaction time for all cases of lime or CaO concentrations in seawater. The yield of magnesium-based compounds is between 20.3% to 26.7% in the first minute of reaction time. The trend appears to be in reciprocation to aragonite yield. For concentrations of CaO of 3 g/L and 5 g/L in seawater, there appeared to be higher concentrations of portlandite or unreacted CaO at 4.8% and 5.6% at the first minute of reaction due to the incomplete CO2 carbonation reaction. Higher calcite levels existing in the first minute for lower Mg/Ca ratios at Ca/Mg of 3 g/L and 5 g/L are also in line with the hypothesis pointed out again, by Bischoff and Fyfe, 1968 [53]; Berner, 1975 [54]; Fernández-Díaz et al., 1996 [55] and De Choudens-Sánchez [56] in which magnesium adsorbed on the surface of calcite seedings to inhibit its further growth. Consequently, we shall see lower calcite yield in the first case [CaO] = 1 g/L or a ratio of Mg/Ca = 1.3353 as more magnesium is able to inhibit calcite crystalline growth. A rule of thumb is proposed here to maximize magnesium compounds yield, which is foremost, to prematurely quench or stop the seawater mineral carbonation at the very beginning of the reaction, which in this case is between 1 min and 5 min to ensure all portlandite reacts with CO2 gas to form calcite seeds subsequently allowing magnesium to adsorb and form complexes on the surface of calcite. Secondly, we kept the ratio of Mg/Ca or concentration of CaO in 33 g/L seawater blends less than 1 g/L or Mg/Ca < 1.3352 to ensure recovery of magnesium is maximized, and finally maintained the temperature at room temperature to ensure that CO2 dissolution is effective to supply the mineral carbonation reaction. Extrapolation of the trending lines in Figure 7a–c pertaining to the magnesium-based compounds yield was performed to estimate the potential yield at 0+ minutes or 30 s reaction time. The trending lines, as extrapolated, showed the highest increase in the magnesium-based compounds precipitation for the case of CaO concentration in seawater at 5 g/L or Mg/Ca = 0.6637, where potential yields were in the range of 50% at 0.5 min or 30 s.
In this case, the operational point for the optimization of magnesium-based compounds was further finetuned to T = 0.5 min, [CaO] = 1 gpl, Temp = 30 °C for the continuous pre-production run targeting some 5000 g of magnesium complex to be synthesized, characterized, and molded into a fire-retardant brick or concrete wall for an accredited fire performance and durability testing.

6. Physico-Chemical Characteristics of Magnesium-Based Compounds Synthesized

The seawater mineral carbonation process was carried out using the in-house patented continuous mineral carbonation reactor with the following reaction parameters finetuned to T = 1 min, [CaO] = 1 gpl, Temperature = 30 °C or at room temperature, with some 5000 g of the magnesium-rich aggregates recovered, filtered, and dried for further characterization. The physico-chemical characteristics of the crystalline precipitate were carried out with the following analytical test methods, which are crystalline microscopy imaging using FESEM, X-Ray diffraction XRD to identify the morphology of the calcium carbonate crystalline precipitate, induction coupled plasma ICP for the presence and concentration of metal ions, and thermal gravimetric analysis TGA to characterize thermal decomposition of the crystalline precipitate.

6.1. FESEM Imaging

The two FESEM images obtained from the model Hitachi FESEM SU8020 are shown in Figure 8 below. With reference to Figure 8, the left image at 15K-times magnification shows an irregular spherical-like cluster shape with rough and textured surface conditions, with an aggregate size of approximately 1 micron to 3 microns. This overall crystallization image may represent early-stage or poorly crystallized aggregates, which possibly formed under premature precipitation. Upon higher resolution, the right image captured at roughly 50K-times magnification shows a well-defined, blocky to cubic-like crystalline shapes with smoother and more regular surface texture at a sub-micron scale. The right image shows a more mature or well-crystallized phase with a resemblance to hydro magnesite or dolomite-type structures.
Previous work by Santoro De Vico et al. (2024) [59], Santos et al. (2023) [60], and De Choudens-Sanchez (2007) [56] also supports the findings above. Furthering the argument, Santoro et al. [59] described the formation of amorphous magnesium carbonate (AMC) as a transient early-stage magnesium carbonate mineralization where the reaction system undergoes a liquid–liquid separation before solid formation, resulting in amorphous and irregular morphologies, while [60] highlighted that hydrated magnesium carbonates tend to be poorly crystalline or amorphous in nature due to kinetic barriers in Mg2+ dehydration. These irregular aggregates are associated with low saturation states and early nucleation, often preceding the formation of more defined polymorphs like calcite or dolomite [61].

6.2. X-Ray Diffractogram XRD Analysis

The hydro magnesite, which has the chemical compound formula of Mg5(CO3)4(OH)2 • 4H2O according to Mineral Data Publishing [62], was subjected to an X-Ray diffraction test provided via Bruker D-8 with library standard matching/interpretations. XRD pattern obtained from the 5000 g products synthesized has a dominant crystallography property matching that of hydro magnesite and calcite, as shown in Figure 9a, according to the interpretation library installed. The hydro magnesite X-Ray diffractograms above are also comparatively similar to the crystallography database of the previous work of researchers such as RRUFF’s American Mineralogy Database [63] in Figure 9b, Akao M et al. (1977) [64], and characterization work on huntite–hydro magnesite geological mineral by Kali Ysilda (2021) [65] as shown Figure 9c. The presence of calcite is due to the quick reaction quenching time at 30 s is in line with the previous results obtained—16.8% of calcite yield detected in Figure 9c when the reaction was quenched at 1 min reaction time.

6.3. Anion Analysis

ICP analysis was carried out on the solid precipitate and spent seawater. Calcium oxide CaO was used as a baseline sample for this study. The baseline results obtained for CaO are tabulated in Table 3 below. It was found that the baseline feedstock CaO has a magnesium content of 4700 ppm and a calcium content of 11,880 ppm. The experiments were conducted at T = 0.5 min, [CaO] = 1 gpl, Temp = 30 °C, and at T = 90 °C.
The solid precipitate synthesized at 30 °C has a spread of magnesium Mg(s) concentration between 7279 ppm and 11,792 ppm and an average of 9962 ppm with respect to the baseline sample lime CaO at 4700 ppm. Hence, magnesium concentration doubled or increased by 211% in the solid precipitate from baseline CaO feedstock. However, at 90 °C reaction temperature, magnesium concentration was still similar to that of our baseline sample, solid lime CaO(s).
The reacted seawater reaction media was vacuum filtered using Whatman Grade 42 filter paper to remove any remnants of solid precipitate several times until the spent seawater appeared to be crystal clear of any opacity. Magnesium Mg2+ ions from seawater were efficiently extracted from the seawater as such from the initial concentration of 1250 ppm and 1322 ppm Mg2+ in raw seawater, the average Mg2+ ion left after the reactions was averaged at 20 ppm, or 98.46% Mg2+ removal from the seawater into solid precipitates.

6.4. Thermal Gravimetric Analysis TGA

TGA on the decomposition properties of the solid precipitate magnesium-rich compounds is shown in Figure 10 below. There are four distinct phases of weight loss identified in the decomposition graph below. The first temperature range is up to 180 °C. Subsequently, the second phase pertaining to the temperature range of 180 °C to 320 °C has an approximate weight loss of 10% and a third distinct phase of a steeper weight loss covering the temperature range of 320 °C to 480 °C, having approximately 20% weight loss. The final fourth phase, from 480 °C to 650 °C, involves another 5% weight loss.
Endothermic degradation is an important criterion to evaluate the effectiveness of a fire-retardant material. Trihydrate alumina and magnesium hydroxide or hydro magnesite are the most well-known mineral fire retardants, as mentioned by Rothon et al. [66]. Hydro magnesite has been proven to decompose endothermically releasing water and CO2 over a temperature range of 220 °C to 550 °C as put out by L.A. Hollingbery et al. (2010) [45] which aligned very much with the results of the TGA curve in this work where the major decomposition of the magnesium compounds also happened at the second and third phase between 180 °C and 490 °C. Haurie et al. (2006) [67] studied the TGA decomposition of synthetic hydro magnesite where a three stage decomposition was observed as such between temperatures of 200–250 °C, there were occurrence of the release of water of crystallization, followed by another release of water from the decomposition of the hydroxides between the temperatures of 380–450 °C and finally, between 510 and 550 °C, there was a release of carbon dioxide from decomposition of the carbonate. This proposed release of CO2 from decomposition of carbonate by Haurie corresponds to the fourth phase of the TGA curve in this work between 480 and 660 °C.
Finally, Inglethorpe et al.’s (2003) [68] findings included that the decomposition of hydro magnesite contains four stages: first endothermic loss of water of crystallization, second endothermic dihydroxylation and formation of amorphous magnesium carbonate; third exothermic crystallization of magnesium carbonate, and finally the endothermic decarbonation of magnesium carbonate, which is very much similar to this work, as represented in Figure 9 above—a four-phase thermal degradation curve.

7. Fire Testing

7.1. Test Methods and Installation Setup

The fire-retardant precipitate from this work was shaped into rectangular blocks for fire testing in accordance with the BS 476-22:1987, as shown in Figure 11 below. Portlandite was added to the ratio of 20:80 to the fire-retardant precipitate to provide minimal support to the blocks subjected to an accredited industrial fire testing method.
The fire resistance test was conducted at the FRIM FIRE TESTING LABORATORY. The procedure undertaken for the fire testing is BS 476-22:1987. Fire Tests on Building Materials and Structures. Methods for the determination of the fire resistance of non-load-bearing elements of construction. The British Standards BS 476-22:1987 states that the fire resistance of the specimen is the time expressed in minutes, to failure under the following main criteria:
  • Integrity—A failure of the test construction to maintain integrity shall be deemed to have occurred when collapse of sustained flaming for more than 10 s on the unexposed face, or one of the following occurs, such as a through gap between 6 and 25 mm exists or develops in the specimen.
  • Insulation—Failure shall be deemed to have occurred when one of the following occurs: the mean temperature of the unexposed surface temperature increases by more than 140 °C above its initial values, and the temperature recorded at any position on the unexposed face is in excess of 180 °C above the initial mean unexposed temperature, or collapse of the mechanical structure of the fire-retardant block.

7.2. BS 476-22:1987—Fire Tests on Building Materials and Structures Results

The actual furnace average temperature versus time is provided in Figure 12a below, with the actual visual of the thermocouples TC1, TC2, and TC3 in Figure 12b, and the construction blueprint diagram in Figure 12c below.
The specimen passed the BS476: Part 22: 1987 after it remained intact for more than 104 min within the integrity specification and remained intact mechanically for more than 96 min, passing the insulation requirement specifications. Overall, with a thin sample thickness of 42 mm, the magnesium complex-rich compound board remained in good condition to perform insulation functions, where it reached an insulation limit of up to 96 min with a maximum temperature of 184 °C. In addition, no integrity failures were detected throughout the 104 min. The real-time unexposed face temperature recorded on the test specimen and observations made throughout the test are provided in Table 4 below.
The testing was stopped at the 104th minute when the early signs of integrity and insulation failure were observed, when a small hole appeared on top of the block sample at 100 min. The visuals of the installation of the magnesium-rich test block are shown as in Figure 13a—before the test, Figure 13b—at 60 min, and Figure 13c—after the test was stopped at 104 min for visual comparison.

8. Conclusions

The hypothesis proposed in this work states that if the process of seawater mineral carbonation is prematurely quenched, magnesium Mg2+ ionic species, which are adsorbed on the calcite lattice formation, will be trapped and therefore recovered in various forms such as magnesium oxides, magnesium hydro magnesite, and magnesium carbonate precipitates were experimentally proven. Hence, a novel method to recover magnesium Mg ions from seawater was successfully experimented and documented as such from the initial 1250 ppm Mg2+ in raw seawater, the average Mg2+ ion left after the reactions was averaged at 20 ppm; or a 98.46% Mg2+ removal from the seawater into the solid precipitates was observed. The continuous seawater mineral carbonation process for the production of magnesium/brucite/huntite products was successfully optimized at T = 0.5 min, [CaO] = 1 gpl, Temp = 30 °C or at room temperature, where the yield of the fire-retardant magnesium-based compounds was 26%. Upon success, a continuous pre-production run targeting a production of some 5000 g of magnesium complex products was successfully conducted, where the products synthesized were characterized and found to have favorable physico-chemical properties for fire retardant/resistant material, as evident from laboratory characterization tests. The solid precipitates were molded into a fire-retardant brick or concrete wall, subjected to an accredited fire performance and durability testing procedure (BS476-22:1987). Encouraging results from the fire resistance testing where the material passed the BS476-22:1987 with the performance criteria such as physical integrity failure, maximum allowable face temperature, and minimum duration up to 104 min.

Author Contributions

Conceptualization, M.G.M.N.; methodology, M.G.M.N.; software, M.G.M.N.; validation, M.G.M.N.; formal analysis, M.G.M.N.; investigation, M.G.M.N.; resources, M.G.M.N.; data curation, M.G.M.N.; writing—original draft preparation, M.G.M.N.; writing—review and editing, M.G.M.N., N.Y.Y., and M.H.b.H.J.; visualization, M.G.M.N.; supervision, N.Y.Y. and M.H.b.H.J.; project administration, N.Y.Y.; funding acquisition, N.Y.Y. All authors have read and agreed to the published version of the manuscript.

Funding

This research was funded by PRSB-UKM R&D Collaboration Project with grant number “KK - 2021- 017” under the title “The Development of Continuous CO2 Mineral Carbonation Process”.During the preparation of this manuscript, I would like to record my sincere acknowledgment of Nasir Darman as the Chief Technology Officer of PETRONAS for endorsing this research grant with UKM.

Data Availability Statement

Restrictions apply to the datasets. The datasets presented in this article are not readily available because the data are part of an ongoing study and are undergoing a market seeding process. Requests to access the datasets should be directed to mghadaffi_mnor@petronas.com.

Acknowledgments

To Faisal Sedaralit, Chief Scientist, for meticulously following the progress of this project from start to end. Last but not least, to my fellow scientists, especially M Syazwan Onn and Ruzilah Sanum, who exhaustively provided their utmost support, energy, and wit day in and day out until the day when we finally achieved the much-sought victory!

Conflicts of Interest

Author Mohammad Ghaddaffi M. Noh is employed by the company PETRONAS Research Sdn.Bhd The remaining authors declare that the research was conducted in the absence of any commercial or financial relationships that could be construed as a potential conflict of interest.

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Figure 1. (a) Chemical compounds that have been successfully synthesized from carbon dioxide gas CO2 from lab scale to commercial scale as compiled by LeClerc (2025) [9] with emphasis given to carbonates. (b) Five types of CO2 reaction based on the number of bonds between the primary carbon and electronegative oxygen (“reduction level”) for the variety of chemical classes.
Figure 1. (a) Chemical compounds that have been successfully synthesized from carbon dioxide gas CO2 from lab scale to commercial scale as compiled by LeClerc (2025) [9] with emphasis given to carbonates. (b) Five types of CO2 reaction based on the number of bonds between the primary carbon and electronegative oxygen (“reduction level”) for the variety of chemical classes.
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Figure 2. Predicted equilibrium solubility product of Mg–calcite. The dashed lines correspond to different reference states for the aqueous ion formation energies [41].
Figure 2. Predicted equilibrium solubility product of Mg–calcite. The dashed lines correspond to different reference states for the aqueous ion formation energies [41].
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Figure 3. The actual image of the continuous reactor skid (left image) used to produce the fire-retardant magnesium-based compounds with components (1) feed tank, (2) positive displacement type pump, (3) 0.8 L continuous mineral carbonation reactor, and (5) backpressure regulator and liquid flow rate valve controller and simplified block diagram (right image) of the laboratory setup of the continuous reactor skid.
Figure 3. The actual image of the continuous reactor skid (left image) used to produce the fire-retardant magnesium-based compounds with components (1) feed tank, (2) positive displacement type pump, (3) 0.8 L continuous mineral carbonation reactor, and (5) backpressure regulator and liquid flow rate valve controller and simplified block diagram (right image) of the laboratory setup of the continuous reactor skid.
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Figure 4. Three-dimensional factor interaction cube to better picture the responses: individual single-factor responses, two-way interaction responses, and three-way factor interaction responses of temperature, CaO concentration, and reaction time to maximize magnesium complex compound yield using seawater as a natural inducer.
Figure 4. Three-dimensional factor interaction cube to better picture the responses: individual single-factor responses, two-way interaction responses, and three-way factor interaction responses of temperature, CaO concentration, and reaction time to maximize magnesium complex compound yield using seawater as a natural inducer.
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Figure 5. Top section, half-normal plot, for the individual factor responses A, B, and C and two-way interaction factor responses AB and BC (top figure), and (bottom section) table comprising the sum of squares and p-value for the individual factor responses A, B, and C; two-way interaction factor responses AB and BC; and final model for the yield of magnesium complex compounds.
Figure 5. Top section, half-normal plot, for the individual factor responses A, B, and C and two-way interaction factor responses AB and BC (top figure), and (bottom section) table comprising the sum of squares and p-value for the individual factor responses A, B, and C; two-way interaction factor responses AB and BC; and final model for the yield of magnesium complex compounds.
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Figure 6. Optimization of the aggregate hydro magnesite, brucite, and huntite yield along the reaction time scale.
Figure 6. Optimization of the aggregate hydro magnesite, brucite, and huntite yield along the reaction time scale.
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Figure 7. (a–c) Results of the optimization of the aggregate hydro magnesite, brucite, and huntite yield along the reaction time scale (min) with reaction temperature at 30 °C and concentration of lime CaO at 1 g/L (or ratio of Mg/Ca = 1.3352), concentration of lime CaO at 3 g/L (or ratio of Mg/Ca = 0.8108) and concentration of lime CaO at 5 g/L or ratio of Mg/Ca = 0.6627.
Figure 7. (a–c) Results of the optimization of the aggregate hydro magnesite, brucite, and huntite yield along the reaction time scale (min) with reaction temperature at 30 °C and concentration of lime CaO at 1 g/L (or ratio of Mg/Ca = 1.3352), concentration of lime CaO at 3 g/L (or ratio of Mg/Ca = 0.8108) and concentration of lime CaO at 5 g/L or ratio of Mg/Ca = 0.6627.
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Figure 8. Images of the synthesized crystalline precipitate with image magnification of 15K times (left) and magnification of 50K times (right) using the continuous mineral carbonation reactor.
Figure 8. Images of the synthesized crystalline precipitate with image magnification of 15K times (left) and magnification of 50K times (right) using the continuous mineral carbonation reactor.
Fire 08 00409 g008
Fire 08 00409 g009
Figure 9. (a) Standard hydro magnesite X-Ray diffraction diagram with reference to RRUFF database [52], (b) XRD pattern obtained for the magnesium compounds precipitation recovered for the case lime CaO concentration in seawater at 5 g/L or Mg/Ca = 0.6637, and (c) X-Ray diffractogram of huntite–hydro magnesite by Kali Ysilda [65].
Figure 9. (a) Standard hydro magnesite X-Ray diffraction diagram with reference to RRUFF database [52], (b) XRD pattern obtained for the magnesium compounds precipitation recovered for the case lime CaO concentration in seawater at 5 g/L or Mg/Ca = 0.6637, and (c) X-Ray diffractogram of huntite–hydro magnesite by Kali Ysilda [65].
Fire 08 00409 g009
Fire 08 00409 g010
Figure 10. Thermal gravimetric analysis TGA of the magnesium compounds synthesized using premature quenching of seawater mineral carbonation techniques via a continuous reactor.
Figure 10. Thermal gravimetric analysis TGA of the magnesium compounds synthesized using premature quenching of seawater mineral carbonation techniques via a continuous reactor.
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Fire 08 00409 g011
Figure 11. Fire-retardant precipitate shaped into rectangular blocks for fire testing in accordance with the BS 476-22:1987 and installed into the fire testing rig at the start of the test, 0 min.
Figure 11. Fire-retardant precipitate shaped into rectangular blocks for fire testing in accordance with the BS 476-22:1987 and installed into the fire testing rig at the start of the test, 0 min.
Fire 08 00409 g011
Fire 08 00409 g012
Figure 12. (a)—Results of the actual furnace temperature time curve; (b) the actual thermocouple TC1, TC2, and TC3 position, and (c) diagram of the thermocouple position during construction.
Figure 12. (a)—Results of the actual furnace temperature time curve; (b) the actual thermocouple TC1, TC2, and TC3 position, and (c) diagram of the thermocouple position during construction.
Fire 08 00409 g012
Fire 08 00409 g013
Figure 13. (a)—visuals of the installation of the magnesium rich test block before the test; (b)—visuals of the block at 60 min; (c) the small holes observed on the top section of the block sample at 100 min shortly before test was halted; and (d) visuals of the test block (face side) after the test was stopped at 104 min.
Figure 13. (a)—visuals of the installation of the magnesium rich test block before the test; (b)—visuals of the block at 60 min; (c) the small holes observed on the top section of the block sample at 100 min shortly before test was halted; and (d) visuals of the test block (face side) after the test was stopped at 104 min.
Fire 08 00409 g013
Table 1. Randomized run for the response “hydro magnesite, brucite, and huntite yield” for the respective factors, temperature, concentration, and reaction time.
Table 1. Randomized run for the response “hydro magnesite, brucite, and huntite yield” for the respective factors, temperature, concentration, and reaction time.
Standard RunRandomized RunFactor AFactor CFactor BResponse Y
Temperature[CaO]Reaction TimeHydro Magnesite/Brucite
Yield
°Cg/Lminutes%
11190601Not detected
22303170.3
34903114.9
414303585.0
52030605Not detected
61930605Not detected
79306011.0
85903114.9
918903515.0
106903114.9
112490605Not detected
1215303583.0
138306012.0
1416903513.6
153303170.1
167306012.0
171090601Not detected
182130605Not detected
1917903580.0
2013303583.0
212290605Not detected
221290601Not detected
232390605Not detected
241303169.9
Table 2. Sum of magnesium complex compound yield Y average responses R1, R2, and R3 from factorial runs pertaining to the respective factors, temperature, concentration, and reaction time, to maximize aragonite yield using seawater as a natural inducer.
Table 2. Sum of magnesium complex compound yield Y average responses R1, R2, and R3 from factorial runs pertaining to the respective factors, temperature, concentration, and reaction time, to maximize aragonite yield using seawater as a natural inducer.
Hydro Magnesite/Brucite
Yield
(Interaction Factors A, B, C)		Run 1Run 2Run 3Sum
1 (L, L, L)70.370.169.9210.0
2 (H, L, L)14.014.914.945.0
3 (L, H, L)2.0215.0
4 (H, H, L)0000
5 (L, L, H)83.085.083.0251.0
6 (H, L, H)13.680.015.0109.0
7 (L, H, H)0000
8 (H, H, H)0000
Table 3. Concentrations of magnesium and calcium ions solid precipitate synthesized at T = 0.5 min, [CaO] = 1 gpl for both temperatures T = 30 °C and T = 90 °C.
Table 3. Concentrations of magnesium and calcium ions solid precipitate synthesized at T = 0.5 min, [CaO] = 1 gpl for both temperatures T = 30 °C and T = 90 °C.
Name of Solid Precipitate SamplesPhaseNaMgCa
ppmppmppm
Calcium Oxide CaO (s)Solid55470011,880
Magnesium compounds at 90 °C_#1Solid244048993620
Magnesium compounds at 90 °C_#2Solid58745133934
Magnesium compounds at 90 °C_#3Solid404010,8153707
Magnesium compounds at 90 °C_#4Solid71872793820
Magnesium compounds at 90 °C_#5solid28311,7921889
Table 4. Unexposed face temperature recorded on the test specimen, and observations made throughout the test until 104 min after the test was discontinued.
Table 4. Unexposed face temperature recorded on the test specimen, and observations made throughout the test until 104 min after the test was discontinued.
Time (min)ThermocouplesMean TemperatureTemperature Rise Above MeanObservation
TC1TC2TC3(°C)Mean TempMax Temp
02728282701Test commences
12728282701
22728282811
32728282811
42728282925
528412933614
6395132411424
7596043542733Observed water moisture appeared at the top surface of the sample at 7 min test
8736760674046
9817267734654
10857674785158
12898179835662
14918582865964
16928683876065
18938885896266
20938886896266
22948987906367
24948988906367
26959090916468
28948988916467
30958990916468
40969090926569All temperature readings still below 100 °C
50989393946771
60999595966972
70999696977072TC1 reached 101 °C at 75 min
801009797987173
901291019710982102
9621114897152125184Small hole appeared on top of the sample at 100 min
10024818497177150221
10427221497194167245Test discontinued
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MDPI and ACS Style

Noh, M.G.M.; Yuhana, N.Y.; Jumali, M.H.b.H. Sustainable Synthesis of Hydro Magnesite Fire Retardants Using Seawater: Characterization, Yield Modeling and Process Optimization. Fire 2025, 8, 409. https://doi.org/10.3390/fire8100409

AMA Style

Noh MGM, Yuhana NY, Jumali MHbH. Sustainable Synthesis of Hydro Magnesite Fire Retardants Using Seawater: Characterization, Yield Modeling and Process Optimization. Fire. 2025; 8(10):409. https://doi.org/10.3390/fire8100409

Chicago/Turabian Style

Noh, Mohammad Ghaddaffi Mohd, Nor Yuliana Yuhana, and Mohammad Hafizuddin bin Hj Jumali. 2025. "Sustainable Synthesis of Hydro Magnesite Fire Retardants Using Seawater: Characterization, Yield Modeling and Process Optimization" Fire 8, no. 10: 409. https://doi.org/10.3390/fire8100409

APA Style

Noh, M. G. M., Yuhana, N. Y., & Jumali, M. H. b. H. (2025). Sustainable Synthesis of Hydro Magnesite Fire Retardants Using Seawater: Characterization, Yield Modeling and Process Optimization. Fire, 8(10), 409. https://doi.org/10.3390/fire8100409

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